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AQA Alevel chemistry
Question | Answer |
---|---|
Name the 3 subatomic particles | Proton, neutron, electron |
What is the relative mass of an electron | 1/2000 or 1/1840 |
What letters do we use to represent mass number and atomic number | Mass number = M and Atomic number = Z |
How is a positive ion formed | By losing electrons |
What ion would be expected to be formed by an element in group 6 | 2- ion |
What is the formula of the sulfate ion | SO42- |
Write the chemical formula of calcium nitrate | Ca(NO3)2 |
Define isotope | Isotopes are atoms of the same element (same number of protons) with different numbers of neutrons |
Calcualte the number of neutrons in C12 and C14 | 6 in C12 and 8 in C14 |
Why do isotopes have the same chemical properties | Same number of electrons |
Describe the atomic model proposed by JJ Thompson | Plum pudding model, sphere of positive charge with electrons dotted around |
Describe the experiment of Rutherford where he discovered the nucleus | Fired alpha particles at gold foil and some deflected and a few bounced back showing small centres of positive charge |
Describe the atomic model proposed by Bohr | Positive nucleus surrounded by electrons in defined shells |
How could you calculate the mass of one atom of oxygen using its Ar of 16 | 16 divided by the avagadro constant (6.022 x 10^23) |
What is meant by relative atomic mass | Mean mass of atoms of an element relative to the mass of 12g of carbon 12 |
Whate are the 5 main stages in mass spectrometry | Vaporisation, Ionisation, acceleration, drift, detection |
Write generic equations to show the ionisation of X in a mass spectrometer by electron impact and proton absorbtion. | e- +X (g) --> X+ (g) + 2e- X + H+ ------> XH+ |
Write a generic equation to show the ionisation of X in a TOF mass spectrometer | H+ +X (g) --> XH+ (g) |
What are the two methods of ionisation that could be used in a mass spectrometer? | electron impact and electrospray |
Describe what happens during electrospray ionisation | Particles gain a proton |
Compare the speed and kinetic energy of a lithium and potassium atom in TOF mass spectrometry | Speed of lithium faster than potassium. Both have same KE |
What is meant by the molecular ion peak | The highest m/z value on the spectrum of a molecule, formed when the molecule does not break up |
How do you work the RMM in a TOF mass spectometer | Subtract one from the mass of the molecular ion |
Why is m/z used in a mass spectrum not just mass | Because atoms can be doubly ionised and would be accelerated more quickly |
How are ions detected in a mass spectrometer? | ions hit the detector and gain electrons, which causes a current to flow |
How can you determine the number of ions hitting the detector at a m/z value? | The size of the current indicates how many ions of that m/z ratio are hitting the detector |
Calculate the Ar of an element sample which is made up of 75% C12 and 25% C14 | 12.5 |
How many electrons can the s, p and d subshells each hold | s = 2, p = 6, d = 10 |
Write the electronic configuration of an element with 20 electrons | 1s2,2s2,2p6,3s2,3p6,4s2 |
Write the electronic configuration of an elements with 24 electrons | 1s2,2s2,2p6,3s2,3p6,4s13d5 |
Which two transition metal elements do not have a full 4s orbital | Chromium and copper |
As a general rule which energy levels do electrons fill first? | Electrons fill orbitals in order from lowest energy to highest energy |
Which subshell fills first 4s or 3d? | 4s |
From which subshell are electrons removed first during ionisation 4s or 3d? | 4s |
Into what orbitals would two electrons in a p subshell go? | Electrons fill orbitals singularly before pairing up |
Fe has electron configuration [Ar] 4s2 3d6. what is the electron configuration of the Fe3+ ion | [Ar] 3d5 |
Define first ionisation energy | The energy required to remove one mole of electrons from one mole of atoms in the gaseous state |
Write an equation to show the 1st ionisation energy of sodium | Na(g) ------> Na+(g) + e- |
Write an equation to show the 4th ionisation energy of calcium | Ca3+(g) ------> Ca4+(g) + e- |
Is ionisation energy exothermic or endothermic | Endothermic - requires energy |
How does nuclear charge affect ionisation energy | Increasing nuclear charge increase ionisation energy as electrons harder to remove |
How does 1st ionisation energy change down the group and give 2 reasons why | Decreases down group. Electrons further from nucleus so less attraction and more shielding so less attraction |
Write an equation to show the third ionisation energy of nitrogen | N2+ --> N3+ + e- |
Why do group 3 elements eg. Al, not fit the trend in increasing first ionisation energy across the period | Electrons removed from a p orbital which is further from the nucleus so less attraction so ionisation energy decreases |
Why do group 6 elements eg. O, not fit the trend in increasing first ionisation energy across the period | Electrons removed from an orbital which contains two electrons which gives some repulsion so easier to remove. So ionisation energy decreases |
What type of bonding exists between metals and non-metals | Ionic bonding |
What happens to electrons in ionic bonding | They are transferred from metal to non-metal |
Describe the nature of an ionic bond | The electrostatic attraction between oppositely charged ions |
Describe the structure of ionic compounds | Giant ionic lattice of positive and negative ions held together by electrostatic forces |
Explain the melting point of ionic compounds | High melting point because strong electrostatic forces within ionic lattice which require a great deal of energy to break |
Explain the electrical conductivity of ionic compounds | Conduct when molten or dissolved because ions can move but do not conduct as a solid because ions held in fixed position in lattice |
Which would you expect to have a higher melting point, NaCl or MgO and why | MgO because ions are double the charge of NaCl so stronger electrostatic forces |
What is the formula of the ionic compound Aluminium Sulfate | Al2(SO4)3 |
What happens to electrons in covalent bonding | Electrons are shared between non-metal atoms |
What is the difference between a single and double covalent bond | Single is 2 shared electrons in bond, double is 4 shared electrons |
Describe the bonding in diamond | Each carbon is bonded to 4 other carbons in a tetrahedral shape |
Describe the structure of graphite | Layers of carbon atoms arranged in hexagonal pattern with weak bonds and delocalised electrons between layers |
Explain why graphite conducts electricity | Delocalised electrons are free to move and carry charge |
Explain the hardness of diamond | Atoms are held in fixed positions |
What is a co-ordinate (dative) bond | Where one atom contributes both electrons in a covalent bond |
Give an example of a molecule with co-ordinate bonding | NH4+, AlCl3NH3, CO |
Why do molecules have different shapes? | Electron pairs repel each other as far apart as possible. |
Which has the greater repulsion effect lone pair of electrons or a covalent bond? | Lone pair |
What is the bond angle in a linear molecule? | 180° |
What is the bond angle in a trigonal planer molecule? | 120° |
What is the bond angle in a tetrahedral molecule ? | 109° |
What shape molecule has 6 bonding electron pairs eg SF6, and what is the bond angle | Octahedral, 90 degrees |
What shape molecule has 5 bonding pairs eg. PCl5 and what are the bond angles | Trigonal bipyramidal, 90 degrees and 120 degrees |
What is the shape and bond angle of CH4. Why | Tetrahedral, 109.5 degree, electron pairs repel each other as far as possible |
How do lone pairs affect bond angles | Reduce bonding pair bond angles (by approx. 2 degrees per lone pair) as they repel more than bonding pairs |
What is the shape and bond angle in water. Why | Bent, 104.5 degree, 4 pairs of electrons including 2 lone pairs |
What is the shape and bond angle in ammonia. Why | Trigonal Pyramidal, 107. degree, 4 pairs of electrons including 1 lone pair |
What is the shape and bond angle in SiCl62- Why | Octahedral, 90. degree, 6 bonding pairs of electrons 0 lone pairs |
Why is the is the bond angle in water less than 180°? | Water contains two bonds and two lone pairs. Lone pairs repel further than bonded pairs of electrons. |
Where in the periodic table are the most electronegative elements | Top right (not including noble gases) |
Which is the most electronegative element | Fluorine |
Which of these bonds will have a dipole? H-C or Cl-C | Cl-C |
What causes molecule to be polar (2 requirements) | significant differences in electronegativity and asymmetrical shape |
Why is CCl₄ not a polar molecule? | Symmetrical molecule so dipoles cancel out. |
Define electronegativity | The ability of an atom to attract electron density (the electron pair) in a covalent bond |
What causes bond polarity | Differences in electronegativity between two atoms bonded together |
Name the 3 types of intermolecular force in order from weakest to strongest | Van der Waals, dipole-dipole and hydrogen bonding |
Which type of intermolecular force exists between diatomic molecules | Van der waals only |
Which types of intermolecular force exist in carbon dioxide | Dipole dipole forces |
Explain how Van der Waals forces arise | Temporary dipoles because of movement of electrons which induce temporary dipoles in neighbouring molecules |
Explain how dipole dipole forces arise | Partial positive charge attracting partial negative charge |
Explain why the boiling point of alkanes increase with increasing chain length | More Van der Waals forces |
Describe the structure of iodine | Molecular lattice of I2 molecules held tougher by Van der Waals forces |
What 3 elements could be bonded to hydrogen to give hydrogen bonding | Oxygen, nitrogen or fluorine |
What is a hydrogen bond? | an electrostatic attraction between a proton in one molecule and an electronegative atom in the other. |
How does hydrogen bonding affect boiling point | Increases boiling point |
Describe the structure of ice | Regular lattice structure of water molecules held together by hydrogen bonding. Less dense than water |
Describe the strucutre of metals | Giant metallic lattice of positive metal ions surrounded by delocalised electrons |
Explain why Mg has a higher melting point than Na | Doubly charged metal ion |
Explain why metals conduct electricity | Delocalised electrons are free to move and carry charge |
Explain why metals have high melting points | The delocalised electrons are strongly attracted to the positive metal ions (strong metallic bonds) which requires a lot of energy to overcome |
What is oxidation in terms of oxygen? | The gain of oxygen |
What is oxidation in terms of electrons? | The loss of electrons |
What does an oxidising agent do? | Donate oxygen/accept electrons (it is reduced) |
What is reduction in terms of oxygen? | The loss of oxygen |
What is reduction in terms of electrons? | The gain of electrons |
What does a reducing agent do? | Accept oxygen/donate electrons (it is oxidised) |
What are redox reactions? | Reactions in which both oxidation and reduction occur |
What is the oxidation state for an element ie Fe/Cl2 | Zero |
What is the oxidation state of a metal in its compound? | The charge on the ion that it forms |
What is the oxidation state of an oxygen atom in a compound? | it is -2 unless it is in a peroxide (H2O2) when it is -1 |
What is the oxidation state of an hydrogen atom in a compound? | it is +1 unless in a hydride (NaI) when it is -1 |
What is the oxidation state of Cl in a compound? | it is -1 unless when bonded to F/O when it can have positive values |
What is disproportionation? | A redox reaction in which the oxidation state of some atoms of a particular element increase in oxidation state and of others decrease in oxidation state |
What is the oxidation state of chromium in the Cr²⁺ ion? | "+2" |
What is the oxidation state of chromium in the CrCl₃? | "+3" |
What is the oxidation number of iron in iron (III) oxide? | "+3" |
What is the oxiddation state of K in K2MnO4? | "+1" |
What is the oxiddation state of O in K2MnO4? | "-2" |
What is the oxiddation state of Mn in K2MnO4? | "+6" |
Write a half equation to show a sodium atom forming a sodium ion | Na → Na+ + e- |
"Write a half equation for the following change: Cr2O72- → Cr3+ | Cr2O7 2-+ 14H+ + 6 e- → 2 Cr3+ + 7 H+ |
Write half equation for MnO4- → Mn2+ and Cl- →Cl2 then combine them | 2x(MnO4- +5e- + 8H+ → Mn2+ + 4H2O) + 5x(2Cl- →Cl2 +2e-) = 2MnO4- + 16H+ + 10Cl- → 2Mn2+ + 8H2O + 5Cl2 |
"Why are elements in Group I (alkali metals) and Group II (alkaline earths) known as s-block elements ? | their highest energy (bonding) electrons are in s orbitals. |
Why does atomic radius increase down group 2? | "the greater the atomic number the more electrons there are; these go into shells increasingly further from the nucleus |
Why does melting point decrease down group 2? | "each atom contributes two electrons to the delocalised cloud, metallic bonding gets weaker due to increased size of ion, Larger ions mean that the electron cloud doesn’t bind them as strongly |
Why does first ionisation energy decrease down group 2? | Despite the increasing nuclear charge the values decrease due to the, extra shielding provided by additional filled inner energy levels |
Why is the second ionisation energy of Magnesium higher than the first ionisation energy? | There are now 12 protons and only 11 electrons. The increased ratio of protons to electrons means that it is harder to pull an electron out |
Why is there such a big jump between the 2nd ionisation energy and the 3rd Ionisation energy of Magnesium? | Because the electron being removed is from a shell nearer the nucleus; there is less shielding |
Write an balanced equation with state symbols for the reaction of Magnesium with steam. | "Mg (s) + H2O (g) → MgO (s) + H2 (g) |
Write an balanced equation with state symbols for the reaction of Barium water | Ba (s) + 2H2O (l) → Ba(OH)2 (aq) + H2 (g) |
What is the trend in solubility of the Group 2 Hydroxides? | They get more soluble down the group. |
Why do the group 2 metals get more reactive as you go down the group? | Reactivity of metals is related to their ionisation energy. The ionisation energies of metals decreases down the group due to increased shielding and increased distance from the nucleus to the outer shell |
What is the trend in solubility of the Group 2 Sulphates? | They get less soluble down the group. |
Why is calcium hydroxide more soluble than magnesium hydroxide and more strongly alkaline? | "Lower charge density of the larger Ca2+ ion means that it doesn’t hold onto the OH¯ ions as strongly. More OH¯ get released into the water. It is more soluble and the solution has a larger pH. |
Describe the chemical test for sulfate ions | Use barium chloride in the presence of dilute hydrochloric acid. This will produce a white precipitate if sulfate ions are present. |
Write an ionic equation with state symbols for the reaction in the sulfate test. | Ba2+ (aq) + SO42– (aq) → BaSO4 (s) |
Write a balanced equation with state symbols for the combustion of Barium in air. | 2Ba (s) + O2 (g) → 2BaO (s) |
Write a balanced equation with state symbols for the reaction of calcium with sulfuric acid | Ca + H2SO4 →CaSO4 + H2 |
Write a balanced equation with state symbols for the reaction of strontium with nitric acid | Sr + 2HNO3 → Sr(NO3)2 + H2 |
Why does the boiling point of Group 7 elements increase down the group? | " increased number of electrons in larger atoms, the van der waals forces between molecules increase so more energy is required to separate the molecules |
Why does the atomic radius of the Group 7 elements increase down the group? | "the greater the atomic number the more electrons there are, these go into shells increasingly further from the nucleus |
Why does the electronegativity of the Group 7 elements decrease down the group? | "There is an increasing number of shells so more shielding and less pull on the outer electrons and an increasing atomic radius so attraction of nucleus for outer electrons drops off as distance increases |
What is the trend in oxidising power of Group 7 elements down the group? | It decreases |
Why does the oxidising power of Group 7 elements decrease down the group? | "the increasing nuclear charge which should attract electrons more is offset by INCREASED SHIELDING and INCREASING ATOMIC RADIUS |
Write a balanced chemical equation with state symbols for the reaction of Chlorine with Sodium Bromide | "Cl2(aq) + 2NaBr(aq) ——> Br2(aq) + 2NaCl(aq) |
Write an balanced ionic equation for the reaction of Bromine with Iodide ions. | "Br2 + 2I¯ ——> I2 + 2Br¯ |
What observation would be made with bromine is reacted with sodium chloride? | "Solution goes from colourless to orange-red |
What observation would be made with chlorine is reacted with sodium iodide? | No visible reaction. |
What observation would be made when solid sodium chloride reacts with concentrated sulfuric acid. | Steamy Fumes |
What a balanced equation for when solid sodium chloride reacts with concentrated sulfuric acid. | "NaCl + H2SO4 → NaHSO4 + HCl |
Write down three observations that would be made when sodium bromide reacts with concentrated sulfuric acid. | steamy fumes, brown fumes, a colourless gas |
Name the three gases produced that would be made when sodium bromide reacts with concentrated sulfuric acid. | Hydrogen Bromide, Bromine and Sulfur Dioxide. |
Write two ionic equations that show the formation of the three gases formed when sodium bromide reacts with concentrated sulfuric acid. | H+ + Br- → HBr and H2SO4 + 2 H+ + 2 Br- → Br2 + SO2 + 2 H2O |
Write down five observations that would be made when sodium iodide reacts with concentrated sulfuric acid. | White fumes, Orange fumes, Colourless gas, Yellow solid, Bad Egg Smell |
Name the three gases produced when sodium Iodide reacts with concentrated sulfuric acid. | "Hydrogen Iodide, Iodine, Sulfur Dioxide, Sulfur, Hydrogen Sulfide |
" | |
Write an ionic equation that shows the formation of Iodine and Sulfur Dioxide when sodium iodide reacts with concentrated sulfuric acid. | H+ + I- → HI |
Write an ionic equation that shows the formation of iodine and sulfur dioxide when sodium iodide reacts with concentrated sulfuric acid. | H2SO4 + 2 H+ + 2 I- → I2 + SO2 + 2 H2O |
Write an ionic equation that shows the formation of iodine and sulfur when sodium iodide reacts with concentrated sulfuric acid. | H2SO4 + 6 H+ + 6 I- → 3 I2 + S + 4 H2O |
Write an ionic equation that shows the formation of iodine and Hydrogen Sulfide when sodium iodide reacts with concentrated sulfuric acid. | H2SO4 + 8 H+ + 8 I- → 4 I2 + H2S + 4 H2O |
Explain why the ability of halide ions to acts as reducing agents increases down the group. | "Down the group it becomes easier to lose an electron because ions are larger & there is more shielding (due to extra electron shell) |
" | |
List the four steps for testing an unknown solution for halide ions | "acidify with dilute nitric acid, add a few drops of silver nitrate solution, treat any precipitate with dilute ammonia solution, if a precipitate still exists, add concentrated ammonia solution |
" | |
What is observation for a positive result to show the presence of chloride ions? | white ppt soluble in dilute ammonia |
What is observation for a positive result to show the presence of bromide ions? | "cream ppt insoluble in dilute ammonia but soluble in conc. |
" | |
What is observation for a positive result to show the presence of iodide ions? | "yellow ppt insoluble in dilute and conc. ammonia solution |
" | |
Write an ionic equation with state symbols for the formation of silver fluoride | "Ag+(aq) + F¯(aq) ——> AgF (s) |
" | |
Write an equation with state symbols for the reactions of chlorine with water | Cl2(g) + H2O(l) <——> HCl(aq) + HOCl(aq) |
Why is the reaction of chlorine with water called a DISPROPORTIONATION reaction? | Because chlorine is simultaneously oxidised and reduced. |
What is the chemical test for chlorine? | "Blue litmus will be turned red then decolourised in chlorine water |
" | |
Write a balanced chemical equation with state symbols for the reaction of Chlorine with cold aqueous sodium hydroxide | 2NaOH(aq) + Cl2(g) —> NaCl(aq) + NaOCl(aq) + H2O(l) |
Explain in terms of the changes in oxidation number why the reaction of Chlorine and cold aqueous sodium hydoride is an example of a DISPROPORTIONATION reaction. | Chlorine changes from 0 to -1 and 0 to +1 in the same reaction. |
What is periodicity | Patterns that recur at regular intervals |
Why do we group elements in the periodic table? | Because they have similar properties |
In what order are the elements of the periodic table arranged? | Increasing atomic number |
Which groups can be classed as s block? | Gp 1+2 |
Which group can be classed as d block? | Transition metals |
Which groups can be classed as p block? | groups 3-7+0 |
Which group can be clased as f block? | lanthanides |
How do you determine whether an element is in an s,p,d/f block? | Depending on which subshell electrons are being added to |
Which energy level are the outer electronsin, in period 3 elements? | 3rd |
Why is sodium described as an s block element? | It is in the s block because the last orbitals to be filled are s orbitals. |
In which block are the transition metals? | d block |
Which groups are in the s block? | groups 1 and 2 |
Define first ionisation energy | The removal of 1 mol of electrons from 1 mol of gaseous atoms. |
Write the equation to show the first ionisation of lithium | Li₍g₎ →Li⁺₍g₎ + e⁻ |
What is the general trend of ionisation energy across period 2? | It increases |
Why does first ionisation energy generally increase across period 2? | Greater attraction between nucleus and outer electrons |
Why are the first ionisation energies of period 3 lower than those of p2 ? | electrons further away and shielded |
Why is there a drop in first ionisation energy between Be and B? | The electrons have started to fill p orbitals which are further away from the nucleus |
Why is there a drop in first ionisation energy between Mg and Al? | The electrons have started to fill p orbitals which are further away from the nucleus |
Why is there a drop in first ionisation energy between N and O? | electrons have started to pair in p orbitals. Electron is repulsed by the other electron making it easier to remove. |
Why is there a drop in first ionisation energy between P and S? | electrons have started to pair in p orbitals. Electron is repulsed by the other electron making it easier to remove. |
Define a dynamic equilibrium. | An reversible reaction in a closed system where both the forward and reverse reaction have the same rate. The concentration of reactants and products remains the same |
What effect does a catalyst have on the equilibrium position? | No effect on equilibrium position. Both Forward and reverse reaction rate increased equally |
How would a temperature increase affect a reversible reaction where the forward reaction is exothermic? | A shift in equilibrium towards the reactants LHS as it reduces the temperature |
What is the expression of Kc for this reaction? A + 2B ⇌ 2C+ D | kc =[A][B]²/[C]²[D] |
What are the units for Kc for this reaction? A + 2B ⇌ 2C+ D | no units |
What are the units for Kc for this reaction? A + 2B ⇌ 2C | mol dm⁻³ |
Kc for a reaction is 0.01 mol dm⁻³. To what side does the equilibrium lie for this reaction? | Towards the LHS more reactants than products in the equilibrium mixture |
Kc for a reaction is 3.23 mol dm⁻³. To what side does the equilibrium lie for this reaction? | Towards the RHS more products than reactants in the equilibrium mixture |
What is an alcohol? | An organic compound with an -OH group |
Draw the structure of the first three alcohols | Methanol, ethanol and propanol |
Name a given alcohol, diol or triol | |
What are the physical properties of alcohols? | Relatively high mp and bp, miscible with water |
Explain the physical properties of alcohols | OH group leads to hydrogen bonding |
What are the two methods for production of ethanol? | Fermentation of glucose, hydration of ethene |
What are the conditions for the production of ethanol from ethene? | Heat, sulphuric acid catalyst, water |
What conditions are necessary for the fermentation of glucose? | Compromise temperature of 35°C, enzymes, sealed vessel |
Give two advantages of production of ethanol from glucose | Low tech required, renewable, potentially carbon neutral |
Give three disadvantages of production of ethanol from glucose | Slow, Low yield, Significant land use, Has to be distilled, Labour intensive |
How is ethanol extracted after fermentation of glucose? | Fractional distillation |
Give two advantages of production of ethanol from ethene | Fast reaction, Pure product, Continuous process, Low on manpower, High yield, 100% alcohol |
Give three disadvantages of production of ethanol from ethene | High technology, Non-renewable source, Expensive equipment |
Write a balanced symbol equation for the complete combustion of a given alcohol | |
What is a dehydration of an alcohol reaction? | One where water is eliminated from an alcohol |
What is the product of a dehydration of an alcohol reaction | An alkene |
What conditions are necessary for dehydration of an alcohol | Conc. phosphoric/sulphuric acid catalyst, heat and reflux |
Draw a mechanism for the acid catalyzed dehydration of a given alcohol | |
What is the advantage of dehydrating alcohols to form alkenes? | Alkenes can be used as chemical feedstock without using crude oil |
What is an oxidation reaction in organic chemistry? | A reaction where oxygen is added or hydrogen is removed from a compound |
What can primary alcohols be oxidized to? | Aldehydes and carboxylic acids |
What can secondary alcohols be oxidized to? | Ketones |
What can tertiary alcohols be oxidized to? | Nothing |
How is oxidation symbolized in an equation? | [O] |
What reagent is used for the oxidation of alcohols? | Acidified potassium dichromate |
What colour change is observed in the oxidation of alcohols? | Acidified potassium dichromate turns green from orange |
Which experimental method is used to convert alcohols to carboxylic acids? | Reflux |
Which experimental method is used to convert alcohols to aldehyde | Distillation |
Which chemical tests can distinguish between aldehydes and ketones? | Tollens's reagent or Fehling's solution |
What is observed when Tollens' reagent is added to aldehydes? | Silver mirror |
What is observed when Fehling's solution is added to aldehydes? | Orange precipitate |
What does collision theory state? | For a reaction to take place particles must collide AND they must collide with enough energy to break the bonds in the reactants |
How does increasing temperature increase the rate of reaction? | This increases the kinetic energy of particles, this increases the frequency and energy of collisions |
How does increasing concentration increase the rate of a reaction? | There are more particles in a given volume, so they are closer together, resulting in more frequent collisions |
How does increasing the pressure of a gas increase the rate of a reaction? | There are more particles in a given volume, so they are closer together, resulting in more frequent collisions |
How does increasing the surface area of a solid increase the rate of reaction? | More particles are exposed and are available to collide, resulting in more frequent collisions |
How does a catalyst increase the rate of a reaction? | Catalysts provide an alternative pathway with a lower activation energy - this means more particles have the minimum energy for successful collisions to occur |
What is activation energy | The minimum amount of energy required to break the bonds in the reactants to allow the reaction to occur |
What is along the x and y axes in a Maxwell Boltzmann distribution graph? | x-energy, y-fraction of particles |
Sketch a typical Maxwell Boltzman distribution curve and indicate the average + most probable energy | |
Why does the Maxwell Boltzmann curve start at zero? | No particles have zero energy |
Why does the right hand side of the curve ever touch the axid at zero | There are no upper limit to the amount of energy some particles may have |
What does the area under a Maxwell Boltzmann curve correlate to? | The total number of particles |
Sketch a typical Maxwell Boltzman distribution curve and another at a higher temperature | |
What does the area under a Maxwell Boltzman curve at and above Eact represent? | The number of particles that have at least the activation energy and above to react |
What happens to the number of particles with Eact and above at higher temperatures? | Increases |
What happens to Eact when a catalyst is used? | The Eact with a catalysts is lower, which results in more particles having the minimum energy in order to react successfully |
What is observed when sodium is added to an alcohol? | Fizzing |
What is observed when potassium dichromate is added to a primary or secondary alcohol? | Turns orange --> green |
What is the test to differentiate between an aldehyde or a ketone? | Fehling's, Benedict's or Tollen's |
What is the test and observation for an alkene? | Bromine water, orange --> colourless |
What are the steps to test for a halogenoalkane? | Add sodium hydroxide, nitric acid, then silver nitrate |
What is the observation when using silver nitrate to test for the different halogenoalkanes? | Chlorine gives white precipitate, bromine cream, iodine yellow |
What is the test and observation for a carboxylic acid? | Add sodium hydrogen carbonate, fizzing |
How can the results from mass spectrometry be used to work out the molecular mass of a compound? | By the mass of the final peak in the spectrum |
How can high resolution mass spectrometry distinguish between molecules with the same molecular mass? | By resolving their masses to a greater/more precise number of decimal places |
What is a wavenumber? | 1/wavelength |
Which regions in an IR spectra will show peaks for a given organic compound | |
What causes global warming? | Absorption of IR waves by bonds in greenhouse gases |
What are the three main greenhouse gases? | Carbon dioxide, methane, water vapour |
What does a negative free energy change tell us? | The process is spontaneous (feasible). |
Define enthalpy change | The heat change at constant pressure |
Define enthalpy of atomisation of an element. | The enthalpy change when one mole of gaseous atoms is formed from the element in its standard state |
Define enthalpy of hydration. | The standard enthalpy change when water molecules surround one mole of gaseous ions. |
Define enthalpy of solution. | The standard enthalpy change when one mole of an ionic solid dissolves completely in sufficient water to form a solution in which the ions are far enough apart not to interact with each other |
Define first electron affinity. | The standard enthalpy change when a mole of gaseous atoms is converted to a mole of gaseous ions each with a single negative charge. |
Define first ionisation energy. | The standard enthalpy change when one mole of gaseous atoms is converted into a mole of gaseous ions each with a single positive charge. |
Define lattice dissociation enthalpy. | The standard enthalpy change when one mole of solid ionic compound is broken up into its free gaseous ions. |
Define lattice formation enthalpy. | The standard enthalpy change when one mole of solid ionic compound is formed from its free gaseous ions. |
Define mean bond enthalpy. | The enthalpy change when one mole of covalent bonds is broken, with all species in the gaseous state, average over a range of different compounds. |
Define second electron affinity. | The standard enthalpy change when a mole of electrons is added to a mole of gaseous ions each with a single negative charge to form ions each with a double negative charge. |
Define second ionisation energy. | The standard enthalpy change when one mole of gaseous unipositive ions is converted into a mole of gaseous ions each with a double positive charge. |
Define standard enthalpy of atomisation of a compound. | The enthalpy change when one mole of a compound in its standard state is converted into its free gaseous atoms. |
Define standard enthalpy of combustion. | The enthalpy change when one mole of a substance is completely burned in oxygen under standard conditions, all reactants and products being in their standard states. |
Define standard enthalpy of formation. | The enthalpy change when one mole of a substance is produced from its constituent elements under standard conditions, all reactants and products being in their standard states. |
Define the Perfect Ionic Model. | A mathematical calculation of the lattice formation enthalpy of a compound, which assumes that the positive and negative ions are perfectly spherical and that there is no covalent character in the compound. |
State Hess's Law | The enthalpy change of a reaction depends only on the initial and final states of the reaction and is independent of the route by which the reaction occurs. |
What are standard conditions? | A standard pressure of 100 kPa and a stated temperature |
What is entropy? | A measure of disorder |
What is the entropy change? | The change in disorder of a system. An increase in disorder produces a positive entropy change. |
What are the Bronsted Lowery definitions of an acid and a base? | An Acid is a proton donor and a Base is a proton acceptor |
What is the conjugate acid of NH3? | NH4+ |
What is the conjugate base of CH3COOH? | CH3COO- |
What is the conjugate acid of HNO3? | H2NO3+ |
Names the Acid Base conjugate pairs in this reaction H3O+ + HSO4- → H2SO4 + H2O | H3O+ Acid H2O Base. HSO4- Base H2SO4 acid. |
Give the formula of Hydroxonium Ion | H3O+ |
Describe the bonding in the Hydroxonium Ion | Covalent and Dative Covalent Bonding |
Define a strong acid | An acid that completely dissociates in aqueous solution |
Define a weak acid | An acid that partially dissociates in aqueous solution. |
What is the general formula for the acid dissociation constant? | Ka= [H+][A-]/ [HA] |
What is the definition of pH? | pH= -log10[H+(aq)] |
What is the pH of a solution with a [H+] of 1.23x10-4 mol dm-3 | 3.91 |
What is the pH of a solution with a [H+] of 7.89 x10-15 mol dm-3 | 14.1 |
Rearrange the definition of pH to make [H+(aq)] the subject | [H+(aq) = 10-pH |
What is the [H+(aq)] of a solution with a pH of 9.2? | 6.31x10-10 moldm-3 |
Calculate the pH of the solution formed by adding of 250 cm3 of water to 50 cm3 of 0.200 mol dm-3 HNO3 | 1.48 |
Calculate the pH of the following solution 100 g dm-3 H2SO4 | -0.31 |
What is the formula of the ionic prodcut of water Kw | Kw = [H+][OH-] |
What is the relationship between pH and pOH? | pH + pOH = 14 |
Calculate the pH of the following solution: 0.20 mol dm-3 Ba(OH)2 | 13.6 |
Calculate the pH of the solution formed by the addition of 100 cm3 of water to 25 cm3 of 0.100 mol dm-3 NaOH | 12.3 |
Write a general equation for the dissociation of a weak acid | "HA(aq) → A¯(aq) + H+(aq) |
" | |
What are the two assumptions one makes when calculating the pH of a weak acid? | "The amount of HA at equilibrium is equal to the amount of HA put into the solution and [H+(aq)]= [A¯(aq)] |
" | |
What does the general formula for the acid dissociation constant become when these two assumptions are applied? | "Ka = [H+(aq)]2/[HA(aq)] |
" | |
Write the fomula you use to calculate the pH of a weak acid. | [H+(aq)] = √ [HA(aq)] Ka |
Calculate the pH of the following weak acid: 0.200 mol dm-3 butanoic acid (Ka = 1.51 x 10-5 mol dm-3) | 2.76 |
Calculate the Ka value for phenylethanoic acid given that a 0.100 mol dm-3 solution has a pH of 2.66 | 4.79 x 10-5 mol dm-3 |
Write the seven steps for the calculation of the pH of a solution formed by the reactions between a weak acid and a strong base if the acid is in excess. | Calculate moles HA (it is still HA and not H+ as it is a weak acid), Calculate moles OH, Calculate moles XS HA, Calculate moles HA left and A- formed, Calculate [HA] leftover and [A-] formed, Use Ka to find [H+], Find pH |
Write the six steps for the calculation of the pH of a solution formed by the reactions between a weak acid and a strong base if the OH- is in excess. | Calculate moles HA (it is still HA and not H+ as it is a weak acid), Calculate moles OH, Calculate moles XS OH, Calculate [OH-], Use Kw to find [H+], Find pH |
Calculate the pH of the solution formed when 50 cm3 of 0.500 mol dm-3 ethanoic acid (pKa = 4.76) is added to 50 cm3 of 0.250 mol dm-3 KOH | 4.76 |
What is a titration? | " a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. |
" | |
What is the equivalence point of a titration? | "the point at which the titration reaction is complete |
" | |
Define what types of substances are suitable to be indicators | They are weak acids whose conjugate base is a different colour to the acid. |
How do you select a suitable indicator for a particular type of titration. | The pH range of the suitable indicator fits within the vertical section of the titration curve. |
"Methyl Orange has a pH range of 3.1-4.4 for what types of titration Is it suitable? | |
" Strong acid, Strong base. Strong acid, Weak Base | |
"Phenolpthalein has a pH range of 8.3-10.0 for what types of titration Is it suitable? | |
" Strong acid, Strong base. Weak acid, Strong Base | |
What is the typical equivalence point of a strong acid strong base titration? | 7 |
What is the typical equivalence point of a strong acid weak base titration? | 5 |
What is the typical equivalence point of a weak acid strong base titration? | 9 |
What is the typical pH range of the vertical section of a strong acid strong base titration curve? | 3 to 10 |
What is the typical pH range of the vertical section of a weak acid strong base titration curve? | 7 to 11 |
What is the typical pH range of the vertical section of a strong acid weak base titration curve? | 3 to 7 |
Explain how one can use a titration curve to calculate Ka of a weak acid. | "Half way to the equivalence point [HA(aq)] = [A¯ (aq) ], so pKa = pH |
" | |
Define what is a buffer solution | "Solutions which resist changes in pH when small quantities of acid or alkali are added. |
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Describe the most common way of making a buffer solution. | "A buffer solution is made from a weak acid and the salt of the weak acid. |
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Describe an additional way of making a buffer solution | "Adding a measured amount of hydroxide ions to a weak acid |
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What are the two assumption we make when calculate the pH of a buffer solution. | "All the A- ions come from the salt, Almost all the HA molecules put into the buffer remain unchanged |
" | |
What does the general formula for the acid dissociation constant become when the assumptions for a buffer solution are applied? | " Ka = [H+(aq)] x [salt]/ [acid] |
" | |
"What happens if an acidic substance is added to a buffer? | |
" "Any rise in [H+(aq)] disturbs the equilibrium. Some A-(aq) ions from the salt react with the extra H+(aq) ions to form HA(aq) and water. A significant fall in pH is prevented. | |
" | |
Calculate the pH of the following buffer solutions made by mixing 50.0 cm3 of 1.00 mol dm-3 methanoic acid (Ka = 1.78 x 10-4 mol dm-3 with 20.0 cm3 of 1.00 mol dm-3 sodium methanoate | 3.35 |
"Calculate the pH of the following buffer solution made by mixing 100 cm3 of 1.00 mol dm-3 ethanoic acid (Ka = 1.78 x 10-4 mol dm-3 ) is mixed with 50.0 cm3 | |
of 0.8000 mol dm-3 sodium hydroxide. | " 3.57 |
Write the balanced redox half equation for the oxidation of iron(ii) to iron(iii) | Fe2+ ——> Fe3+ + e¯ |
"Write the balanced redox half equation for the MnO4¯ being reduced to Mn2+ in acidic solution | |
" MnO4¯ + 5e¯ + 8H+ ———> Mn2+ + 4H2O | |
"Write the balanced redox half equation for Cr2O72- being reduced to Cr3+ in acidic solution | |
" Cr2O72- + 6e¯ + 14H+ ——> 2Cr3+ + 7H2O | |
"Fill in the gaps in this phrase: Each electrode / electrolyte combination has its own half-reaction which sets up a ---------- ---------- because different cells have a different tendency to donate or accept ---------. | |
" "Each electrode / electrolyte combination has its own half-reaction which sets up a potential difference because different cells have a different tendency to donate or accept electrons. | |
" | |
On what variables does Electrode Potential depend? | "TEMPERATURE, PRESSURE OF ANY GASES, SOLUTION CONCENTRATION |
" | |
Define Standard Electrode Potential | "The definition of standard electrode potential E° of a half cell is the voltage of that half cell relative to a standard hydrogen electrode when all solutions have a concentration of 1 moldm-3 all gases are at 1 atmosphere pressure and the temperature |
" | |
What is the Standard Electrode Potential of a standard hydrogen half cell defined as? | 0.00V |
What is the purpose of a salt bridge between two half cells? | it completes the circuit by allowing the transfer of ions. |
How does one make a salt bridge? | Soak filter papers in a saturated solution of potassium nitrate. |
By convention how are all half equations written for each standard electrode potential? | All equations are written as reduction processes |
If two half equations are combined how do you know what the overall reaction will be? | "A species with a higher E° value oxidise (reverses) one with a lower value |
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Write an acronym to remember how to write a combination of cells in standard cell notation. | Reduced Oxidised Oxidised Reduced. ROOR |
Write a formula to work out Ecell for a cell written in cell notation. | "Ecell= E°RHS-E°LHS |
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Write a formula to work out Ecell for a redox reaction. | "Ecell= E° half equation for the Reduction reaction - E° half equation for the oxidation half equation. |
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Write down the 7 rules for writing a cell in cell notation. | "ROOR, Add State Symbols, Species in the same state in the same half cell are separated by a comma, Species in different states are separated by a solid vertical line, If there is not solid species in the half cell you must add a Platinum Electrode, If th |
Water molecules are optional | |
" | |
Write the oxidation and reduction reactions occurring, and then write the overall reaction occurring in each cell for the Ag/Zn electrode pair. | "Zn(s) → Zn2+ (aq) + 2e- and Ag+(aq) + e- → Ag(s) |
so overall: Zn(s) + 2Ag+ (aq) → Zn2+ (aq) + 2Ag(s)" | |
Calculate the standard electrode potential of the Na+/Na electrode given that when it was joined to the standard hydrogen electrode, the cell emf was –2.71 volts. | -2.71 |
Calculate the emf of a cell with the standard AgCl/Ag electrode (Eº = +0.22 V) as the left hand electrode and the Fe2+/Fe (Eº = –0.44 V) electrode as the right hand one. | emf = –0.44 – +0.22 = –0.66 V |
Calculate Eºcell of the following cells using the Eº values from a data sheet: Ni2+/Ni Sn4+/Sn | 0.44 |
Calculate Eºcell of the following cells using the Eº values from a data sheet: 2I-/I2, Ag+/Ag | 0.26 |
"How can E° values be used to predict the feasibility of redox and cell reactions? | |
" "In theory ANY REDOX REACTION WITH A POSITIVE ECELL VALUE WILL WORK, A half equation with a more positive E° value reverse a less positive one | |
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Give two reasons why a redox reaction that is predicted to be feasible might not occur | Its activation energy might be too high, it might not be occurring under standard conditions. |
How do the redox half equations that are set up in a rechargeable battery when the battery is being recharged compare to when it is being charged | They are reversed by an external current. |
What are the pros and cons of non rechargeable batteries. | Pros cheap, small Cons waste issues due to very limited lifetime |
What are the pros and cons of rechargeable batteries. | "pros: less waste, cheaper in long run, Cons time to recharge, still some waste issues because lifetime is still limited |
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In general, what is a d-block element? | An element whose outermost electron is in a d orbital |
In general, what is a ligand substitution reaction? | A reaction in which one or more ligands is/are replaced by one or more different ligands in a complex |
In general, what is a transition element? | An element which has at least one stable ion with a partially filled d-orbital |
List the four typical properties of a transition metal | "Variable oxidation states The formation of coloured aqueous ions. The formation of complex ions Their use as Heterogeneous Catalysts |
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What is the electronic structure of copper? | "1s2 2s2 2p6 3s2 3p6 4s1 3d10 |
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Why is zinc not a transition metal? | It does not form a stable ion which has unfilled d orbitals |
Explain in terms of ionisation energy why Transition metals have a variety of oxidation states | "Transition metals have a large number of oxidation states because when removing d electrons there is never a big jump from one ionisation energy to the next. |
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List the five colours of the five oxidation states of vanadium | " yellow VO2+, blue VO2+ ions, green V3+ ions, V2+ |
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Define bidentate ligand | A ligand which can donate two lone pairs of electrons to form two coordinate bonds with the central metal ion |
Define complex | A central metal ion surrounded by ligands |
Define coordinate (dative covalent) bond. | An electrostatic force of attraction between the shared pair of electrons and the nuclei of the atoms involved where both electrons are from the same atom. |
Define coordination number. | the number of dative covalent bonds formed to a central metal ion in a complex |
Define ligand. | A molecule or ion which can donate one or more lone pairs of electrons to form coordinate (date covalent) bond(s) with a central metal ion in a complex |
Define monodentate ligand. | A ligand which can donate one lone pair of electrons to form one coordinate bond with the central metal ion |
Define multidentate ligand | A ligand which can donate multiple lone pairs of electrons to form multiple coordinate bonds with the central metal ion |
Describe the chelate effect and how it affects the stability of ligands | "When we replace ligands with those that form more co-ordinate bonds, reaction is feasible (driven by increase in entropy) |
Reverse reaction is NOT feasible due to large +ve ΔG caused by big decrease in entropy | |
" | |
"Name this complex ion [Cr(H20)6]3+ | |
" "hexaaquachromium(III) ion | |
" | |
"Name this complex ion [Cu(CN)6]4- | |
" Hexacyanocuprate(II) | |
"Name this complex ion [Co(H2O)4Cl2]+ | |
" Tetraaquadichlorocabalt(III) | |
Write an equations for the reaction of chloride ions with Hexaaquacopper(II) | [Cu(H2O)6]2+ + 4Cl- → [CuCl4]2- + 6H2O |
Write an equation for the reaction of excess aqueous ammonia with hexaaqucopper(II) ions | [Cu(H2O)6]2+ + 2NH3 → [Cu(H2O)4(OH)2] + 2NH4+ (hydrolysis) |
Write an equations for the reaction of a few drops of aqueous ammonia with hexaaqucopper(II) ions | [Cu(H2O)6]2+ + 4NH3 → [Cu(H2O)2(NH3)4]2+ + 4H2O |
What are the two possible shapes for complex ions with a coordination number of 4 | Tetrahedral and square planar |
Explain the origin of the colour of transition metal aqua ions such as [Cr(H2O)6]3+. | "d – orbitals/subshell/energy level split (in energy by (ligands), Electron transitions/jumps from lower to higher energy level, Absorbs light in visible region/reference to white light |
" | |
Explain why [Zn(H2O)6]2+ ions have no colour | " d-(sub) shell / orbitals are full / 3d10 arrangement of electrons (1), No jumps of d-electrons /no d-d transitions |
" | |
How does a catalyst work? | "substance that speeds up reaction without being used up, They provide an alternative mechanism with lower activation energy. |
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Define a heterogenenous catalyst | "Heterogeneous catalyst is in different phase to the reactants. |
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Define a homogenenous catalyst | Homogeneous catalyst is in the same phase as the reactants. |
Describe the three stages of a heterogeneneous catalyst work? | Reactants adsorbed onto surface (onto active sites), Reaction takes place, Products are desorbed (leave the surface). |
Write two equations that show the action of Fe2+ as a homoegeneous catalyst for this reaction 2 I- + S2O82- -------> I2 + 2 SO42- | "2 Fe2+ + S2O82- → 2 Fe3+ + 2 SO42- , 2 Fe3+ + 2 I- → 2 Fe2+ + I2 |
" | |
Define Lewis acid and Lewis base | Lewis acid- electron pair acceptor, Lewis base-electron pair donor |
Give the formula of metal-aqua ions formed when iron(II), copper(II), iron (III) and aluminium ions in aqueous solution. | [Fe(H2O)6]2+, [Cu(H2O)6]2+,, [Fe(H2O)6]3+,[Al(H2O)6]3+, |
Explain why [Fe(H2O)6]3+, is more acidic than [Fe(H2O)6]2+, | Fe3+ is more charge dense than Fe2+ , because it is smaller and has a higher charge. This means it has a bigger polarising effect on its water ligands. This weakens the 0—H bond in the water ligand and makes it more likely to hydrolyse in the presence of |
Describe what you would observe and give an equation for the reaction between [Fe(H20)6]2+ and a few drops of sodium hydroxide. | Green solution forms a green precipitate [Fe(H2O)6]2+ + 20H- ---> [Fe(H20)4(OH)2 + 2H2O |
"Describe what you would observe and give an equation for the reaction between [Fe(H20)6]2+ and a few drops of ammonia | |
. | " Green solution forms a green precipitate [Fe(H20)6]2+ + 2NH3 ------> [Fe(H20)4(OH)2] + 2NH4+ |
State and explain what happens to the transition metal containing product of the above reaction if the reaction mixture is left to stand in air. | The green precipitate turns brown as the iron (II) is oxidised to iron (Ill) by the oxygen in air |
Describe what you would observe and give an equation for the reaction between [Fe(H20)6]2+ and a few drops of sodium carbonate. | Green solution forms a green precipitate [Fe(H20)6]2+ + CO32 - ------.> FeCO3 + 6H20 |
Describe what you would observe and give an equation for the reaction between [Cu(H20)6]2+ and a few drops of sodium hydroxide. | Blue solution forms a Blue precipitate [Cu(H2O)6]2+ + 20H- ---> [Cu(H20)4(OH)2] + 2H2O |
"Describe what you would observe and give an equation for the reaction between [Cu(H20)6]2+ and a few drops of ammonia | |
. | " Blue solution forms a Blue precipitate [Cu(H20)6]2+ + 2NH3 ------> [Cu(H20)4(OH)2] + 2NH4+ |
"Describe what you would observe and give an equation for the reaction between [Cu(H20)4(OH)2] and excess aqueous ammonia | |
. | " Blue precipitate to dark blue solution [Cu(H2O)4(OH)2] + 4NH3 → [Cu(H2O)2(NH3)4]2+ + 2H2O + 2OH- |
Describe what you would observe and give an equation for the reaction between [Cu(H20)6]2+ and a few drops of sodium carbonate. | Blue solution forms a Blue precipitate [Cu(H20)6]2+ + CO32 - ------.> CuCO3 + 6H20 |
Describe what you would observe and give an equation for the reaction between [Al(H20)6]3+ and a few drops of sodium hydroxide. | Colourless solution forms a white precipitate [Al(H2O)6]3+ + 30H- ---> [Al(H20)3(OH)3] + 3H2O |
"Describe what you would observe and give an equation for the reaction between [Al(H20)3(OH)3 and excess Sodium Hydroxide | |
. | " White precipitate re-dissolved to form a colourless solution [Al(H2O)3(OH)3] + OH- → [Al(H2O)2(OH)4]- + H2O |
Describe what you would observe and give an equation for the reaction between [Al(H20)6]33+ and a few drops of sodium carbonate. | Colourless solution forms a white precipitate effervescence 2[Al(H20)6]3+ + 3CO32 - ------.> 2[Al(H20)3(OH)3] + 3CO2 + H2O |
"Describe what you would observe and give an equation for the reaction between [Al(H20)6]3+ and a few drops of ammonia | |
. | " Colourless solution forms a White precipitate [Al(H20)6]3+ + 3NH3 ------> [Al(H20)3(OH)3] + 3NH4+ |
Describe what you would observe and give an equation for the reaction between [Fe(H20)6]3+ and a few drops of sodium hydroxide. | Orange solution forms a brown precipitate [Fe(H2O)6]3+ + 30H- ---> [Fe(H20)3(OH)3] + 3H2O |
"Describe what you would observe and give an equation for the reaction between [Fe(H20)6]3+ and a few drops of ammonia | |
. | " Orange solution forms a brown precipitate [Fe(H20)6]3+ + 3NH3 ------> [Fe(H20)3(OH)3] + 3NH4+ |
Describe what you would observe and give an equation for the reaction between [Fe(H20)6]33+ and a few drops of sodium carbonate. | Orange solution forms a brown precipitate effervescence 2[Fe(H20)6]3+ + 3CO32 - ------.> 2[Fe(H20)3(OH)3] + 3CO2 + H2O |
Write equations to demonstrate the amphoteric nature of [Al(H20)3(OH)3] | [Al(H20)3(OH)3] + OH- ------> [Al(OH)4]- + 3H2O [Al(H20)3(OH)3] + H3O+ ------> [Al(H2O)6]3+ + 3H2O |
Write the balanced equation for sodium reacting with water including state symbols | 2Na(s) + H2O(aq) → 2NaOH(aq) + H2(g) |
Which product determines the pH of the solution formed when sodium reacts with water. | NaOH(aq) it is alkaline (pH13-14) |
What pH is the solution formed after sodium reacts with water? | pH 13-14 |
Why does sodium melt into a ball when it reacts with water | The reaction is highly exothermic |
Describe and explain the difference in reactivity between sodium and magnesium (4) | Na is more reactive, Mg contains stronger metallic bonds, due to the metal ions having a higher charge and being smaller and because there are more delocalised electrons |
Why is the reaction of magnesium with water much slower than sodium | The metallic bonding in magensium is stronger causing the activation energy of the reaction to be high |
Write a balanced equation for magensium reacting with steam, including state symbols | Mg(s) + H2O(g) → MgO(s) + H2(g) |
Write a balanced equation for sodium reacting with oxygen, including state symbols | Na(s) + O2(g) → Na2O(s) |
Write a balanced equation for magnesium reacting with oxygen, including state symbols | 2Mg(s) + O2(g) → 2MgO(s) |
Write a balanced equation for aluminium reacting with oxygen, including state symbols | 4Al(s) + 3O2(g) → 2Al2O3(s) |
Write a balanced equation for silicon reacting with oxygen, including state symbols | Si(s) + O2(g) → SiO2(s) |
Write a balanced equation for phosphorus reacting with oxygen, including state symbols | 4P(s) + 5O2(g) → P4O10(s) |
Write a balanced equation for sulfur reacting with oxygen to form sulfur dioxide, including state symbols | S(s) + O2(g) → SO2(g) |
Write a balanced equation for sulfur reacting with oxygen to form sulfur trioxide, including state symbols | 2SO2(s) + O2(g) → 2SO3(g) |
Describe the structure and bonding in sodium, magensium and aluminium oxide | giant ionic lattice |
Describe and explain the difference in the mpt of sodium and magnesium oxide | mpt of MgO is higher because the metal ion is smaller with a higher charge resulting in greater electrostatic forces between oppositely charged ions |
Describe and explain the difference in the mpt of magensium and aluminium oxide | MgO is higher because Al2O3 contains covalent character, due to the very small and highly charged metal ion which distorts the electron density of the O2- ion |
Describe the structure and bonding in silicon dioxide | giant covalent structure |
Explain why the mpt of silicon dioxide is very high | due to giant covalent structure, strong covalent bonds need to be broken which requires a lot of energy |
Describe the structure and bonding of phosphorus and sulfur oxides | simple covalent molecules |
Which has a higher mpt P4O10 or SO2. Explain your answer | P4O10, as it is a larger molecule resulting in stronger van der waals forces between molecules |
Which period 3 oxides are basic? | sodium and magensium oxide |
Write a balanced equation for sodium oxide reacting with water, including state symbols | Na2O(s) + 2H2O(l) → 2NaOH(aq) |
Write a balanced equation for magnesium oxide reacting with water, including state symbols | MgO(s) + 2H2O(l) → Mg(OH)2(aq) |
What is the pH og Mg(OH)2 | About 8-10 |
Why does NaO dissolve to make an alkaline solution? | Na breaks up into its ions, the O2- ion takes an H+ from water forming 2 OH- ions |
Why is the pH of sodium hydroxide higher than magnesium hydroxide? | MgO is only slightly soluble, due to a higher lattice enthalpy. The concentration of O2- ions in solution are low therefore the OH- concentration is low, making the solution only slightly alkaline |
What is the pH of the solution when Al2O3 is mixed with water and why | 7, as it is insoluble due to a high lattice enthalpy |
What is the pH of the solution when sulfur dioxide is mixed with water and why | 7, as it is insoluble due giant covalent structure which needs a large amount of energy to break |
Which period 3 oxides are acidic? | phosphorus and sulfur oxides |
Write a balanced equation for phosphorus oxide reacting with water, including state symbols | P4O10(s) + 6H2O(l) → 4H3PO4(aq) |
Write a balanced equation for sulfur dioxide reacting with water, including state symbols | SO2(g) + H2O(l) → H2SO3(aq) |
Write a balanced equation for sulfur trioxide reacting with water, including state symbols | SO3(g) + H2O(l) → H2SO4(aq) |
Write a balanced equation including state symbols to show how sulfuric (VI) acid dissociates into its ions in solution | H2SO4(aq) → 2H+(aq) + SO42-(aq) |
Write a balanced equation including state symbols to show how sulfuric (IV) acid dissociates into its ions in solution | H2SO3 (aq)→ H+(aq) + HSO3-(aq) |
Write a balanced equation including state symbols to show how phosphoric acid dissociates into its ions in solution | H3PO4(aq)→ H+(aq)+ H2PO4-(aq) |
Which period 3 oxides react with acids? | sodium and magensium oxides |
Write a balanced equation including state symbols to show how sodium oxide reacts with hydrochloric acid | Na2O(s) + 2HCl(aq) → 2NaCl(aq) + H2O(l) |
Write a balanced equation including state symbols to show how sodium oxide reacts with sulfuric acid | Na2O(s) + H2SO4(aq) → Na2SO4(aq) + H2O(l) |
Write a balanced equation including state symbols to show how magnesium oxide reacts with hydrochloric acid | MgO(s) +2HCl(aq) → MgCl2(aq) + H2O(l) |
Write a balanced equation including state symbols to show how megnesium oxide reacts with sulfuric acid | MgO(s) + H2SO4(aq) → MgSO4(aq) + H2O(l) |
Which period 3 oxides react with bases? | phosphorus and sulfur oxides |
Write a balanced equation including state symbols to show how silicon dioxide reacts with sodium hydroxide | SiO2(s) + 2NaOH(aq) → Na2SiO3(aq) + H2O(l) |
Write a balanced equation including state symbols to show how phosphorus oxide reacts with sodium hydroxide | P4O10(s) + 12NaOH(aq) → 4Na3PO4(aq) + 6H2O(l) |
Write a balanced equation including state symbols to show how sulfur dioxide reacts with sodium hydroxide | SO2(g) + 2NaOH(aq) → Na2SO3(aq) + H2O(l) |
Write a balanced equation including state symbols to show how sulfur trioxide reacts with sodium hydroxide | SO3(g) + 2NaOH(aq) → Na2SO4(aq) + H2O(l) |
What does amphoteric mean? | Something that reacts with both acids and bases |
Which period 3 oxide is amphoteric and why? | Aluminium oxide, because it has both ionic and covalent character |
Write a balanced equation including state symbols to show how aluminium oxide reacts with sodium hydroxide | Al2O3 + 6NaOH (aq) 3H2O(l) → 2NaAl(OH)4(aq) |
Write a balanced equation including state symbols to show how aluminium oxide reacts with sulphuric acid | Al2O3 + 3H2SO4(aq) → Al2(SO4)3 (aq) + 3H2O(l) |
Write a balanced equation including state symbols to show how aluminium oxide reacts with hydrochloric acid | Al2O3 + 6HCl(aq) → 2AlCl3(aq) + 3H2O(l) |
What is an empirical formula? | The simplest whole number ratio of atoms of each element in a compound |
What is a molecular formula? | The formula that shows the actual number of atoms of each element in a molecule |
What is a general formula? | A generalised formula which enables you to work out the molecular formula of a compound within a homologous series |
What is a structural or shorthand formula? | A formula showing the arrangement of atoms in a substance without showing the bonds. E.g. CH₃CH₂CH₃ |
What is a displayed formula? | A formula showing the arrangement of atoms in a substance and all the bonds E.g. |
What is a skeletal formula? | A simplified version of a displayed formula where straight lines represent carbon-carbon bonds e.g. |
What is a hydrocarbon? | A compound made of atoms of hydrogen and carbon only |
What is an alkane? | A hydrocarbon with only single bonds. |
What is chemical nomenclature? | The process of naming different organic compounds |
Name a given branched alkane | Assign parent chain, assign branches a number and add as prefixes |
What is isomerism? | When molecules have the same atoms but they are arranged differently in space |
What is structural isomerism? | Where isomers have the same molecular formula but a different structure |
What are the three types of structural isomerism? | Chain, position and functional |
What is a saturated hydrocarbon? | A hydrocarbon with only single bonds |
What is crude oil? | A mixture of mainly alkanes that can be separated by fractional distillation |
What is fractional distillation? | The process of separating the alkanes in crude oil by use of their different melting points |
What is cracking? | The process of breaking C-C bonds in alkanes to make shorter molecules |
What are the two types of cracking? | Thermal cracking and catalytic cracking |
How is thermal cracking carried out? | At high pressure and temperature |
What is thermal cracking used for? | Producing a high percentage of alkenes |
How is catalytic cracking carried out? | Slight pressure, high temperature, zeolite catalyst |
What is catalytic cracking used for? | Production of motor fuels and aromatic hydrocarbons |
Why is cracking carried out? | The products of cracking are in greater demand than the long chain alkanes |
What is combustion? | The exothermic reaction of a fuel with oxygen |
What is complete combustion? | A combustion reaction where there is enough oxygen to ensure products are purely carbon dioxide and water |
What is incomplete combustion? | A combustion reaction where the lack of oxygen results in products other than carbon dioxide and water |
What can the products of incomplete combustion be? | Carbon monoxide, carbon |
Which pollutants are produced by the internal combustion engine? | Nitrogen oxides, carbon monoxide, carbon, unburned hydrocarbons |
What is the purpose of a catalytic convertor? | To remove gaseous pollutants produced by internal combustion engines |
How is sulphur dioxide released from combustion engines? | Through the combustion of hydrocarbons containing sulphur impurities |
What is a flue gas? | The gas released from large industrial plants and power plants |
How can sulphur dioxide be removed from flue gases? | By using calcium oxide or calcium carbonate to neutralise it |
What is a free radical? | A reactive species containing an unpaired electron |
How are free radicals signaled? | By use of a ∙ e.g.. Cl∙ |
How are chlorine free radicals produced? | Exposure to UV light |
What is free radical substitution? | Where an atom bonded to a carbon is substituted for a free radical atom |
Which steps are involved in a free radical substitution? | Initiation, propagation, termination |
What is an initiation step? | The generation of a free radical |
Give an example of an initiation step | Cl₂ → 2Cl∙ |
What is a propagation? | Intermediate steps in a free radical reaction which cause the regeneration of the original free radical |
Give an example of a propagation | CH₄ + Cl∙ → CH₃∙ + HCl, CH₃∙ + Cl₂ → CH₃Cl + Cl∙ |
What is a termination? | A reaction involving two free radicals and resulting in no free radicals |
Give an example of a termination | CH₃∙ + Cl∙ → CH₃Cl |
Why do free radical substitutions result in impure products | Because there are many possible side reactions e.g. further substitution with chlorine or alkyl radical-radical reactions |
What is a halogenoalkane? | A saturated hydrocarbon with a halogen bonded to at least one carbon |
Name a given halogenoalkane | |
What explains halogenoalkanes increased reactivity relative to alkanes? | The polar carbon-halogen bond |
Why do halogenoalkanes contain polar bonds? | Due to the difference in electronegativity between carbon and halogens |
What kind of reactions do halogenoalkanes undergo? | Nucleophilic substitution |
What is a nucleophile | A substance with a lone pair of electrons |
Which three nucleophiles do halogenoalkanes react with? | OH⁻, CN⁻ and NH₃ |
What reagents and conditions are necessary for hydroxyl substitutions of halogenoalkanes? | Reflux in aqueous solution of NaOH |
What reagents and conditions are necessary for nitrile (cyanide) substitution of halogenoalkanes? | Aqueous, alcoholic KCN |
What is the importance of nucleophilic substitution with cyanide ions? | Increases the length of the carbon chain |
What reagents and conditions are necessary for ammonia substitution of halogenoalkanes? | Reflux in aqueous, alcoholic excess ammonia |
Why is excess ammonia required in its nucleophilic substitution with a halogenoalkane? | To prevent further substitutions |
Draw a nucleophilic substitution mechanism for a given halogenoalkane reacting with a given nucelophile | |
How is the rate of reaction dependent on the enthalpy of the C-X bond? | The greater the enthalpy of the C-X bond, the greater its strength, the slower the reaction |
What is the trend in C-X reaction rates | Increases down the group from C-F to C-I |
What is an elimination reaction? | A reaction where a halogenoalkane becomes an alkene |
Which elements are "eliminated" in an elimination reaction | H and X |
What conditions favour elimination reactions? | Secondary or tertiary halogenoalkane, hot ethanolic KOH |
Draw the mechanism for the elimination of a given halogenoalkane with OH⁻ | |
What is ozone? | O₃ |
Where is ozone found and why is it beneficial? | Upper atmosphere, absorbs harmful UV radiation |
What is a CFC? | A chlorofluorocarbon: a compound containing C-Cl and C-F bonds |
What were CFCs used for? | Refrigerants and solvents |
What happens to CFCs when exposed to UV? | C-Cl bond breaks and generates Cl∙ |
How do chlorine radicals interact with ozone? | Catalyse its decomposition into O₂ thus forming a "hole" in the ozone layer |
Why are CFCs now banned for use? | Due to research results from different groups within the scientific community |
Give two equations showing the decomposition of ozone with chlorine radicals | Cl∙ + O₃ → ClO∙ + O₂ , ClO∙ + O₃ → 2O₂ + Cl∙ |
What is the overall equation for the decomposition of ozone? | 2O₃ → 3O₂ |
What is an alkene? | A hydrocarbon with a carbon-carbon double bond |
Why are alkenes described as unsaturated? | They have a carbon-carbon double bond |
Describe the electron density of the double bond | It is a region of high electron density |
Draw the shape of the carbon-carbon double bond orbitals | |
What type of isomerism can alkenes exhibit? | Stereoisomerism (E-Z or cis-trans) |
What physical feature of alkenes causes their stereoisomerism? | That the double bond cannot rotate |
Name a given alkene | |
What two type of reactions do alkenes typically undergo? | Electrophilic addition or polymerisation |
Which three reagents do alkenes typically react with? | HBr, H₂SO₄ and Br₂ |
What type of reagents do alkenes react with? | Electrophiles |
What are the products of an electrophilic addition reaction with Br₂? | A dibromoalkane |
Draw a reaction mechanism for a given alkene with Br₂ | |
Why is bromine water used as a test for alkenes? | The Br₂ reacts with the alkene so its reddish-brown colour is decolourised |
What are the products of an addition reaction with HBr? | A bromoalkane |
What is a carbocation intermediate? | An intermediate molecule in a reaction which has a positive charge on a carbon atom |
From a given carbocation, state whether it is primary, secondary or tertiary | |
Describe the variation in stabilities of the different orders of carbocation | Tertiary > secondary > primary |
How does stability of carbocations effect the products of an addition reaction with HBr? | More stable carbocation is formed resulting in a major and minor product |
Draw a mechanism for an addition reaction of an alkene and HBr. Draw and label the major and minor products. | |
What is the product of a reaction with an alkene and concentrated H₂SO₄? | An alcohol |
Draw a mechanism for the reaction between a given alkene and concentrated sulphuric acid | |
Why can sulphuric acid be considered a catalyst in hydration of alkenes? | It is regenerated by the end of the reaction. |
What are addition polymers formed from? | Alkenes and substituted alkenes |
How are addition polymers named? | Poly(name of original alkene) |
Draw the polymer for a given substituted alkene | |
Draw the monomer for a given section of polymer | |
What are the chemical properties of addition polymers? | Generally unreactive |
Why are addition polymers unreactive? | Strong C-C and C-H bonds |
What is PVC? | Poly(chloroethene) |
What is a plasticiser? | A small molecule that can get inbetween polymer chains resulting in them being able to slide over one another more easily |
What is the relative atomic mass of an atom? | The average mass of an atom of an element (taking into account its naturally occurring isotopes)4 relative to 1/12 the relative atomic mass of an atom of carbon-12 |
What is the relative molecular mass of a molecule? | The mass of a molecule compared to 1/12 the relative atomic mass of an atom of carbon-12 |
How are the number of moles of a substance established from its mass? | Moles = mass/Mr |
How many moles are in Xg of Y? | X/Mr of Y |
How many moles are there in 4.00 kg of CuO | 50.3 |
How many moles are there in 39.0 g of Al(OH)3 | 0.5 |
How many moles are there in 1 tonne of NaCl | 17100 |
How many moles are there in 20.0 mg of Cu(NO3)2 | 0.000107 |
What is the mass of 0.200 moles of Al2O3 | 20.4g |
What is the mass of 0.00200 moles of (NH4)2SO4 | 0.264 g |
What is the mass of 0.300 moles of Na2CO3.10H2O | 85.8 g |
Calculate the number of moles of Al3+ ions in 5.10 g of Al2O3. | 0.1 |
It was found that 1.00 g of vitamin C contains 0.00568 moles of Vitamin C molecules. Calculate the Mr of vitamin C. | 176 |
0.8500 g of hexanone, C6H12O, is converted into its 2,4-dinitrophenylhyrazone during its analysis. After isolation and purification, 2.1180 g of product C12H18N4O4 are obtained. Calculate the percentage yield. | 88.40% |
What is the mass of X moles of Y? | X x Mr of Y |
X moles of a substance has a mass of Y. What is its Mr? | Y/X |
In what molar ratio do hydrochloric acid with calcium hydroxide react? | One to two |
In what molar ratio do nitric acid with ammonia react? | One to one |
What mass of aluminium reacts with 258 mg of chlorine? 2Al + 3Cl2 → 2AlCl3 | 0.0654g |
In what molar ratio do sulfuric acid with barium hydroxide react? | One to One |
In what molar ratio do nitric acid and pottasium carbonate react? | Two to One |
"Deduce the limiting reagent and calculate what mass of magnesium oxide is formed when 486 mg of | |
magnesium reacts with 240 mg of oxygen 2Mg + O2 → 2MgO | " Oxygen. 0.605g |
What is the atom economy of a reaction? | A ratio of the mass of desired product to the mass of all products |
How is atom economy calculated? | (Mr or mass of desired product/Mr or mass of all products) x 100 |
What is the atom economy to make tungsten in this reaction: WO3 + 3H2 → W + 3H2O | 77.30% |
"Calculate the mass of aluminium oxide that would be formed when 2.70 g of aluminium reacts with | |
2.56 g of oxygen.4Al + 3O2 → 2Al2O3 | " 5.10g |
At 273 K and 101000 Pa, 6.319 g of a gas occupies 2.00 dm3 . Calculate the relative molecular mass of the gas. | 71 |
Calculate the atom economy when titanium is extracted from titanium chloride. TiCl4 + 2Mg → Ti + 2MgCl2 | 20.10% |
What standard units are used for concentration of a solution? | mol/dm3 |
How are cm³ converted to dm³? | Divide by 1000 |
What are concordant titres? | Titres within 0.1cm³ of each other |
What colour is phenolphthalein in acid and alkali? | Colourless in acid, purple in alkali |
State the ideal gas equation | PV = nRT |
What is the empirical formula of a substance? | The formula that represents the simplest whole number ratio of the atoms of each element present in the compound. |
What is the molecular formula? | The actual number of atoms of each element in one molecule of a compound |
What is an ionic equation? | A simplified equation showing the ions present in a reaction |
What is a spectator ion? | An ion that does not take part in a reaction |
Which substances produce ions in ionic equations? | Aqueous solutions of ionic compounds |
"Write an ionic equation, including state symbols, for the reaction of aqueous ammonia with hydrochloric acid | |
" H+(aq) + NH3(aq) → NH4+(aq) | |
Write an ionic equation, including state symbols for the precipitation of lead(II) bromide when aqueous lead(II) nitrate is mixed with aqueous sodium bromide | Pb2+(aq) + 2Br–(aq) → PbBr2(s) |
"Write an ionic equation, including state symbols, for the reaction of potassim carbonate solution with nitric acid | |
" H+(aq) + CO32-(aq) → H2O + CO2(g) | |
Write an ionic equation, including state symbols for the precipitation of lead(II) iodide when aqueous lead(II) nitrate is mixed with aqueous potassium iodide | Pb2+(aq) + 2Br–(aq) → PbBr2(s) |
What volume of carbon dioxide gas, measured at 800 K and 100 kPa, is formed when 1.00 kg of propane is burned in a good supply of oxygen? C3H8 + 5 O2 → 3 CO2 + 4 H2O | 4.53 m3 |
0.140 moles of a gas has a volume of 2.00 dm3 at a pressure of 90.0 kPa. Calculate the temperature of the gas | 155 K |
"What volume of oxygen is required to burn the following gases, and what volume of carbon dioxide is produced? 1 dm3 of methane CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l) | |
" "O2 2 dm3 | |
, CO2 1 dm3" | |
If 4 dm3 of hydrogen sulphide is burned in 10 dm3 of oxygen, what is the final volume of the mixture (give the volume of each gas at the end)? 2 H2S(g) + 3 O2(g) → 2 H2O(g) + 2 SO2(g) | 4 dm3 O2, 4 dm3 H2O, 4 dm3 SO2, total 12 dm3 gas |
Calculate the number of moles in the following. 2 dm3 of 0.05 mol dm-3 HCl | 0.1 |
"Calculate the number of moles in the following. 50 litres of 5 mol dm-3 H2SO4 | |
" 250 | |
Calculate the number of moles in the following. 10 cm3 of 0.25 mol dm-3 KOH | 0.0025 |
Calculate the concentration of the following in both mol dm-3 and g dm-3 0.400 moles of HCl in 2.00 litres of solution | 0.2 mol dm-3 , 7.3 g dm-3 |
Calculate the concentration of the following in both mol dm-3 and g dm-3 12.5 moles of H2SO4 in 5.00 dm3 of solution | 2.5 mol dm-3 245.3 g dm-3 |
Calculate the concentration of the following in both mol dm-3 and g dm-3 1.05 g of NaOH in 500 cm3 of solution | 0.0512 mol dm-3 , 2.10 g dm-3 |
25.0 cm3 of a solution of sodium hydroxide required 18.8 cm3 of 0.0500 mol dm-3 H2SO4. H2SO4 + 2 NaOH → Na2SO4 + 2 H2O Find the concentration of the sodium hydroxide solution in mol dm-3 | 0.0752 mol dm-3 |
What volume of 5.00 mol dm-3 HCl is required to neutralise 20.0 kg of CaCO3? 2 HCl + CaCO3 → CaCl2 + H2O + CO2 | 79.9 dm3 |
Limestone is mainly calcium carbonate. A student wanted to find what percentage of some limestone was calcium carbonate. A 1.00 g sample of limestone is allowed to react with 100 cm3 of 0.200 mol dm-3 HCl. The excess acid required 24.8 cm3 of 0.100 mol dm | 87.70% |
An impure sample of barium hydroxide of mass 1.6524 g was allowed to react with 100 cm3 of 0.200 mol dm-3 hydrochloric acid. When the excess acid was titrated against 0.228 mol dm-3 sodium hydroxide in a back titration, 10.9 cm3 of sodium hydroxide soluti | 90.80% |
Write the empirical formula of each of the following substances. a) C2H6 b) P2O3 c) SO2 d) C6H12 | a CH3 b P2O3 c SO2 d CH2 |
" The empirical formula and relative molecular mass of some simple molecular compounds are shown below. Work out the molecular formula of each one. b) C2H5 Mr = 58 c) CH2 Mr = 70 e) CH Mr = 78 f) CH2 Mr = 42 | |
" b) C4H10 c) C5H10 e) C6H6 f C3H6 | |
50.0 g of a compound contains 22.4 g of potassium, 9.2 g of sulphur, and the rest oxygen. Calculate the empirical formula of the compound. | K2SO4 |
A compound contains 40.0 g of carbon, 6.7 g of hydrogen and 53.5 g of oxygen. It has a relative molecular formula of 60. Find both the empirical and the molecular formula of the compound. | CH2O, C2H4O2 |
A compound contains 59.4% carbon, 10.9% hydrogen, 13.9% nitrogen and 15.8% oxygen, by mass. Find the empirical formula of the compound. | C5H11NO |
25.0 cm3 of 0.0400 mol dm-3 sodium hydroxide solution reacted with 20.75 cm3 of sulphuric acid in a titration. Find the concentration of the sulphuric acid. | 0.0241 mol dm-3 |
What is enthalpy? | A measure of the heat energy of a substance |
What is enthalpy change ΔH? | The change in heat energy of a substance at constant pressure |
What is the unit of ΔH? | KJmol-1 |
What is ΔHθ? | The change in heat energy under standard conditions - 100kPa + 298K |
Are breaking bonds exo/endothermic? Explain your answer | Endo - energy is taken in to break the bonds |
Are making bonds exo/endothermic? Explain your answer | Exo- energy is given out when bonds are made |
Draw an energy level diagram for an exothermic reaction | |
Draw an energy level diagram for an endothermic reaction | |
How do you calculate the overall energy change of a reaction? | Energy released when bonds are made in the products- energy needed to break bonds in the reactants |
In an exothermic reaction is the enthalpy change positive or negative? | negative |
Why is the enthalpy change in an exothermic reaction negative? | The products have less energy than the reactants because energy is lost to the surroundings |
In an endothermic reaction in the enthalpy change positive or negative? | positive |
Why is the enthalpy change in an endothermic reaction positive? | The products have more energy than the reactants because energy is taken in from the surroundings |
What is the standard enthalpy of formation ∆Hf? | The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions |
Write an equation to show the standard enthalpy of formation of ethanol | 2C(s) + 3H2(g) + 1/2 O2(g) → C2H5OH(l) |
What is the standard enthalpy of combustion ∆Hc? | The enthalpy change when one mole of a compound is completely burned in oxygen in standard conditions |
Write an equation to show the standard enthalpy of combustion of ethene. | C2H4(g) + 3O2(g) → 2CO2(g) + 2H2O(l) |
What is the mean bond enthalpy (bond dissociation enthalpy)? | The average energy needed to break a certain type of bonds in a range of compounds |
Why are bond enthalpies always positive? | energy is required to break bonds - endothermic |
Why might the bond enthalpy calculated using Hess's law be different to the mean bond enthalpy from a data book? | mean bond enthalpies aren't exact, they are averaged over a range of compounds |
How can you use bond enthalpies tocalculate the enthalpy change for a reaction? | (∑ bond enthalpies of products)-(∑ bond enthalpies of reactants) |
What is the specific heat energy of a substance © | The amount of energy needed to raise the temperature 1g of substance by 1K |
What is the unit of specific heat capacity | Jg-1K-1 |
Describe the experiment you would do to calculate the enthapy change of a combustion reaction? | Calorimetry-burn a known amount of reactant and record the ∆T of known mass of water |
What is the equation you can use to calculate enthalpy change? | q=mc∆T |
Describe each of the components in the q=mc∆T with units | q = heat lost/gained in J, m-mass of solution in g/water in calorimetry, ∆T change in temperature of the water Kelvin(K), c=specific heat capacity of solution/water JK-1K-1 |
Describe the experiment you would do to calculate the enthalpy change of an exo/endo reaction | carry out the reaction in a polystyrene cup, measure the temperature every minute starting before mixing, plot time vs temp on a graph to obtain ∆T on mixing by extrapolating |
Why is extrapolation used to find an accurate ∆T for an exo/endo reaction? | to allow for a heat loss from the polystyrene cup |
Suggest why the experimental ∆Hc of ethanol is lower than the data book value | Some heat is lost to the surroundings and is not all transferred to the water |
Suggest what could be done to reduce the heat loss in a calorimetry experiment | use a heat shield to prevent the heat lost from the burning substance |
What is Hess's law? | The overall enthalpy change for a reaction is the same independent of the route taken |
Draw a thermochemical cycle that you would draw to find out ∆H for the following reaction if given ∆Hf data, C2H2 + H2 → C2H6 | |
Draw a thermochemical cycle that you would draw to find out ∆H for the following reaction if given ∆Hc data, C2H2 + H2 → C2H6 | |
What is the ∆Hf of O2 and why? | Zero, because the ∆Hf of elements in their standard states are zero |
What are the two main types of isomerism? | Structural isomerism and stereoisomerism |
What are the three types of structural isomerism? | Chain isomerism, position isomerism, functional group isomerism |
What are the two types of stereoisomerism? | Geometric isomerism and optical isomerism |
Which organic molecules demonstrate optical isomerism? | Ones containing chiral carbons |
How can chiral carbons be identified? | Carbon atoms with four different groups surrounding them |
State whether a given molecule exhibits optical isomerism | |
What are enantiomers? | Molecules with the same structural formula but exhibiting optical isomerism |
How do enantiomers differ in terms of their mirror images? | They are non-superimposable upon their mirror images |
How do enantiomers differ in terms of their effect on plane polarized light? | They rotate it in opposite directions |
What is a racemic mixture? | A 50:50 mixture of two enantiomers |
Explain the effect of a racemic mixture on plane polarized light | One enantiomer rotates the light clockwise and the other anti-clockwise, resulting in no net rotation. |
Which homologous series are considered "carbonyls"? | Aldehydes and ketones |
How can aldehydes and ketones be distinguished? | Fehling's solution or Tollen's reagent |
Draw a named carbonyl compound | |
Name a given carbonyl compound | |
How can aldehydes be formed? | Oxidation of a primary alcohol or reduction of a carboxylic acid |
How can ketones be formed? | Oxidation of a secondary alcohol |
Name the reagent and conditions for the reduction of carbonyl compounds | NaBH₄, aqueous |
How is the reductant represented in equations? | [H] |
What is the name of the mechanism for the reduction of carbonyls? | Nucleophilic addition |
Draw a mechanism for the reduction of a given carbonyl | |
Under which conditions are hydroxynitriles formed? | Addition of KCN followed by dilute acid to a carbonyl |
Draw a mechanism for the reaction of KCN with a given carbonyl | |
Explain how nucleophilic additions of carbonyls with KCN can result in a racemic mixture | Carbonyl is planar, KCN can attack from above or below with equal probability, 50:50 chance of forming each enantiomer |
Which hazards are associated with using KCN? | It is highly toxic |
Write an overall equation for the formation of a given hydroxynitrile with HCN | |
Why do aldehydes and ketones have higher boiling points than alkanes with the same number of carbon atoms? | The carbonyl bond means that there are permanent dipole dipole interactions between the molecules which are not present between the alkane molecule |
Why do aldehydes and ketones have lower boiling points than alcohols with the same number of carbon atoms? | There is hydrogen bonding between the alcohol molecules. This is a stronger intermolecular force than the dipole dipole attractions between the carbonyl compounds |
Why are small carbonyl compounds miscible in water? | The formation of hydrogen bonds from the water to the oxygen of the carbonyl compound. |
Why is the carbonyl bond polar? | The difference in electronegativity between carbon and oxygen |
What is the difference in structure between an aldehyde and ketone? | "ALDEHYDES have at least one H attached to the carbonyl group. Ketones have two carbons attached to the carbonyl group. |
" | |
Name this molecule CH3CH2COCH2CH3 | "pentan-3-one |
" | |
Draw the structural formulae of this molecule phenylethanone | |
Explain in detail the technique recrystallization. | Dissolve the solute in the minimum volume of hot solvent, cool the solution to allow the crystals to form, remove the crystals by vacuum filtration, dry the resulting crystals in an oven. |
What measurement of a substance is the best measure of its purity | Melting Point |
Give the name of a suitable reducing agent for the reduction of a carbonyl compound to an alcohol. | Lithium Tetrahyridoaluminate, Sodium Tetrahydridoborate |
What is the nucleophile produced by NaBH4 and LiAlH4? | H- |
"Draw the mechanism for the reduction of butanal to butanol | |
" | |
"Why is Propanoyl chloride more easily reduced by Hydride ions than Propanal. | |
" "Chlorine are more electronegative than carbon, Atoms attached to the carbonyl carbon that are more electronegative than carbon make it more delta positive (1) | |
More delta positive the carbonyl carbon is the more reactive it is, towards nucleophiles such as H-, Therefore carbonyl carbon is most delta positive in propanoyl chloride (1) and propanoyl chloride is the most reactive with H-. | |
" | |
What are the reagents for the nucleophilic addition of cyanide ions to aldehydes or ketone? | "hydrogen cyanide - HCN in the presence of KCN |
" | |
What are the conditions for the nucleophilic addition of cyanide ions to aldehydes or ketone? | "reflux in alkaline solution |
" | |
What is the structural formula of the product of the reaction of these two molecules CH3CHO + HCN | "CH3CH(OH)CN |
" | |
Draw out the mechanism for the reaction between these two molecules CH3CHO + HCN | |
What is the empirical formula of this Hydrocarbon Carbon = 92% Carbon (Atomic mass = 12) , Hydrogen = 8 % (Atomic mass = 1) | CH |
What is the molecular formula of this Hydrocarbon Carbon = 92% Carbon (Atomic mass = 12) Hydrogen = 8 % (Atomic mass = 1), Relative molecular mass = 78 | C6H6 |
What is the unit of rate of reaction? | mol dm-3 s-1 |
Write a rate equation for a reaction that is first order with respect to reactant A and 2nd order with reactant B. | r = k [A] [B]2 |
For a reaction that is second order with respect to reactant A, what will happen to the rate of reaction if the concentration of A is doubled? | It will be quadrupled. |
Explain what the order of a particular reactant is in terms of the particles of that reactant in the mechanism of the reaction | The order of a particular reactant is the number of particles of that reactant that participate before or during the rate determining step. |
What is the order of a particular reactant in the rate equation of a particular reactant | "The power to which a concentration is raised in the rate equation |
" | |
What is the overall order of a chemical reaction. | "The sum of all the individual orders in the rate equation. |
" | |
If a reaction is zero order with respect to a reactant Z what is the effect on the rate of a reaction of increasing the the concentration of Z | No effect. |
If a reaction is zero order with respect to a reactant Z where must Z be in the mechanism of the reaction? | After the Rate Determining Step |
What is the unit of the rate constant k in this rate equation? rate = k [A] | s-1 |
What is the unit of the rate constant k in this rate equation? rate = k [P] [C] | mol-1 dm3 s-1 |
What is the unit of the rate constant k in this rate equation?rate = k [D]2 [B] | "mol-2 dm6 s-1 |
" | |
How does a catalyst increase the rate of a reactant? | "Catalysts provide an alternative reaction pathway with a lower Activation Energy. Decreasing the Activation Energy means that more particles will have sufficient energy to overcome the energy barrier and react |
" | |
The Arrhenius equation is lnK = lnA- Ea/RT. Describe how a graph that could be plotted from a number of values of K(rate constant) and T (Temperature in K) to determine the value of activation energy Ea | lnK should be plotted on the y axis, 1/T on the x axis then the gradient is -Ea/R |
Draw the structure that Kekule proposed for Benzene | |
What chemical test shows that Benzene molecules do not contain any double bonds? | It does not decolourise bromine water. |
Describe the shape and structure of a Benzene molecule in full | hexagon, planar, all carbon carbon bonds the same length. |
How do the carbon carbon bond lengths in a benzene molecule prove it cannot be made up of alternate double and single carbon carbon bonds. | All the carbon carbon bonds are the same length. |
The enthalpy of hydrogenation of cyclohexene is -120KJ/Mol, The enthalpy of hydrogenation of benzene is -208KJ/Mol. By how many KJ/Mol is Benzene more stable than you would expect | 152KJ/mol |
" What word describes how the electrons are delocalized (spread) across the whole Benzene molecule? | |
" Conjugation. | |
How does valence bond theory describe the bonding in Benzene? | "unhybridized p-orbitals which overlap forming a cloud of electron density above and below the molecule. |
" | |
Give four pieces of evidence that Benzene does not contain the Kekule structure? | The carbon -carbon bond lengths are all the same, Benzene does not undergo electrophilic addition reactions, Benzene molecules are 152KJ/Mol more stable than the Kekule structure. |
Describe the flame that is produced when Benzene is burnt in air | Yellow and Smokey |
Explain why Benzene burns with such a yellow and smoky flame. | Benzene has a very high carbon hydrogen bond ratio. |
What are the conditions required to convert Benzene to Cyclohexane by reaction with Hydrogen | Rayney Nickel Catalyst at 150oC |
Why do electrophiles that react with benzene have to be stronger than those that react with alkenes | "Because the mechanism involves an initial disruption to the ring. The delocalized electrons in the benzene ring make it more stable. |
" | |
What are the required reagents and conditions for the nitration of benzene. | "conc. nitric acid and conc. sulphuric acid (catalyst), reflux at 55°C |
" | |
Draw out the mechanism for the nitration of benzene | |
Write the equation for the formation of the nitrate electrophile when nitric acid reacts with concentrated sulfuric acid. | "2H2SO4 + HNO3 <----> 2HSO4¯ + H3O+ + NO2+ |
" | |
Write the equation for the formation of the chlorine electrophile which is formed by the reactions of FeCl3 with Cl2 | "Cl2 + FeCl3 ------------> FeCl4¯ + Cl+ |
" | |
What are the conditions for the chlorination of benzene | "reflux in the presence of a halogen carrier e.g. Fe, FeCl3, AlCl3 |
" | |
Draw the structure of the product of the reaction between Benzene and fuming sulfuric acid. | |
What are the reagents for the Friedel Crafts Alkylation of benzene | Reagents a halogenoalkane (RX) and anhydrous aluminium chloride AlCl3 |
What are the conditions for the Friedel Crafts Alkylation of benzene | "Conditions reflux ; dry inert solvent (ether) |
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Write an equation for the Friedel Crafts Alkylation of benzene with chloroethane. | "C6H6 + C2H5Cl ———> C6H5C2H5 + HCl |
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Draw out the mechanism for the Friedel Crafts alklyation by chloroethane. | |
What are the reagents for the Friedel Crafts Acylation of benzene | "Reagents an acyl chloride (RCOX) and anhydrous aluminium chloride AlCl3 |
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What are the conditions for the Friedel Crafts Acylation of benzene | "Conditions reflux 50°C; dry inert solvent (ether) |
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Write an equation for the Friedel Crafts Acylation of benzene by ethanoyl chloride | "C6H6 + CH3COCl ———> C6H5COCH3 + HCl |
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What does DNA stand for? | Deoxyribonucleic acid |
What is an amino acid? | A biological molecule with two functional groups |
What is DNA? | A large molecule containing genetic information |
What is formed from different amino acids combined in one chain? | Proteins |
What is the structure of DNA? | Two polymer chains made from nucleotides in the form of a double helix |
Which functional groups do amino acids have? | NH₂ and -COOH |
Draw the structure of lysine at low pH | |
Draw the structure of alanine when it is in the presence of concentrated sodium hydroxide. | |
Draw the zwitterion of alanine. | |
Draw the structure of the organic product when valine reacts with methanol in the presence of concentrated sulfuric acid. | |
Draw the structure of the organic product when alanine reacts with ethanoyl chloride. | |
Draw the structure of the repeating unit formed when valine polymerises. | |
Draw the structure of the organic product when an excess of bromoethane reacts with alanine. | |
Describe the composition of proteins. | Proteins are sequences of amino acids joined by peptide bonds |
List three levels of protein structure. | Primary, secondary (α-helix and β-pleated sheet), tertiary |
Draw the structure of the tripeptide formed from ala-val-ala | |
Draw the structure of amino acids formed when the given peptide is hydrolysed. | |
Explain how hydrogen bonds, ionic bonds and sulfur-sulfur maintain the structure of proteins. | "Hydrogen bonding tends to occur between the partially negatively charged oxygen of -C=O of one amino acid and the partially positively charged hydrogen of -N—H on a different amino acid. Ionic bonds can form between side chain groups of amino acids. This |
Sulfur—sulfur bonds occur most commonly when cysteine is present in the protein structure. This amino acid has a -CH2SH side group that is capable of reacting together under suitable oxidising conditions. This creates a disulfide bridge." | |
Name and describe a simple method to separate a mixture of amino acids. | Thin-layer chromatography can be used to separate amino acids. This uses a plate coated with a thin layer of silica (SiO2) that acts as the stationary phase. The mixture dissolves in a suitable mobile phase and, as it moves up the plate, the amino acids |
Give an equation used to calculate Rf values | Rf = distance moved by the spot / distance moved by the solvent |
Amino acids are invisible. Give two methods of making them visible on the TLC plate. | The plate can be sprayed with a developing agent such as ninhydrin in order to make the spots visible. Alternatively, ultraviolet light can be shone on the plate in order to make the spots visible. |
Describe the structure of an enzyme. | Enzymes are proteins that are folded into complex shapes that allow smaller molecules to fit in them. |
Paclitaxel is a commonly used anticancer drug. Explain, in detail, how it can Enzyme-substrate complex interact with an enzyme. | "Enzymes contain an active site where the chemical reaction happens. This active site is specific to a |
particular substrate. When attached, the reacting molecule is held in the right orientation to react. It | |
is important that the reacting molecule is able to bond temporarily to the enzyme through | |
intermolecular forces." | |
How can drugs be used as enzyme inhibitors? | If a drug that mimics the shape of the substrate is introduced to the body it can block the active site of the enzyme to the substrate. This process is called enzyme inhibition. |
Outline the role of computers in developing enzyme inhibitors | Computers can be used to model the shapes of proteins and enzymes and, therefore, can allow medicinal chemists to model the effect of introducing a particular enzyme inhibitor to that protein. |
Why is a racemic mixture of a drug often less potent than an optically pure sample? | Active sites are stereospecific, which means that they are capable of binding to only one enantiomeric form of a substrate or drug. Therefore, 50% of a racemic mixture would not bind properly to the active site. |
Describe the structure of a nucleotide. | Nucleotides are the monomers that make up DNA. They are comprised of a phosphate ion bonded to 2-deoxyribose, which is then bonded to one of the four DNA bases (cytosine, guanine, adenine and thymine). |
Describe how nucleotides join together to form a single strand of DNA | Strands of DNA are formed when nucleotides polymerise. Covalent bonds between the phosphate group and the 2-deoxyribose of different nucleotides. The polymer takes the form of a sugar-phosphate-sugar-phosphate chain with different bases attached to the su |
How do the single strands of DNA form together in a double helix? | Hydrogen bonds form between bases on each chain, which causes them to join together in a double helix |
State the two pairs of DNA bases that are able to hydrogen bond to each other. | Adenine and thymine, cytosine and guanine |
Draw the structure of cisplatin | |
State the most common medicinal use of cisplatin. | Cisplatin is most commonly used as an anticancer drug |
Fill in the gaps for this explanation of how cisplatin prevents DNA, Cisplatin works by bonding to guanine bases in ........... This occurs by a ...... ............ reaction which cisplatin prevents DNA replication in cells. with a new bond forming be | Cisplatin works by bonding to guanine bases in DNA. This occurs by a ligand replacement reaction which cisplatin prevents DNA replication in cells. with a new bond forming between platinum and a nitrogen atom on guanine. Chloride ions are displaced bec |
Outline some of the major side effects of cisplatin. | " Nausea, weakened immune system, mouth ulcers, loss of appetite, loss of fertility |
" | |
Why can drugs such as cisplatin have such adverse side effects? | Anticancer drugs are carried by blood around the body and have an impact on every cell they come into contact with. They have the greatest effect on rapidly dividing cells such as hair, reproductive cells and cells that line the stomach |
How do oncologists mitigate against these side effects in their patients? | Dosages spread out at intervals of three weeks, regular checks of white blood cell count to lower risk of developing fever, steroids to lower nauseous feelings |
Why might patients consider not taking cisplatin or other anticancer drugs? | Although cisplatin can increase the remaining time a patient has to live, it can severely lower the quality of life of the patient due to the side effects. Every patient, along with their oncologist, has to weigh up the balance between the quantity an |
Complete the following statement - All nuclei with the same chemical environment have the same __________ ___________ | Chemical Shift |
How many peaks (not including splitting) are found in the 1H NMR spectrum of butanoic acid? | 4 |
How many peaks (not including splitting) are found in the 1H NMR spectrum of ethyl ethanoate? | 3 |
How many peaks (not including splitting) are found in the 1H NMR spectrum of methyl propanoate? | 3 |
How many peaks (not including splitting) are found in the 1H NMR spectrum of propyl methanoate? | 4 |
How many peaks are found in the 13C NMR spectrum of butanoic acid? | 4 |
How many peaks are found in the 13C NMR spectrum of ethyl ethanoate? | 4 |
How many peaks are found in the 13C NMR spectrum of methyl propanoate? | 4 |
How many peaks are found in the 13C NMR spectrum of propyl methanoate? | 4 |
In 1H NMR Spectroscopy what does the integration data tell us? | Relative number of equivalent H atoms in the molecule |
In 1H NMR Spectroscopy what does the splittin pattern tell us? | The number of adjacent, non-equivalent protons according to the n + 1 rule |
In NMR Spectroscopy what does the number of peaks relate to? | The number of different environments |
In NMR Spectroscopy what the chemical shift tell us? | The molecular environment |
What is the n + 1 rule? | It gives the number of peaks into which the signal of the neighbouring environment will be split |
What is TMS? | Tetramethylsilane. Its used as the standard for NMR Spectroscopy and its peak is set to zero on the chemical shift scale |
Why are the advantages of TMS? | Volatile, inert, it has signals that are well to the right of the majority of the peaks in organic chemistry and it has one environment |
Why is deuterated trichloromethane used as the solvent in 1H NMR spectroscopy (instead of undeuterated trichloromethane)? | Deuterium is not NMR active and therefore doesn’t show up on the spectrum - it is still used in 13C but the peak can removed by the computer software |
What is the only way that the value of Kc or Kp for a particular equilibrium can be changed | A change in temperature |
"What effect does changes to any of the following have on the value of the equilibrium constant: concentration of reactants or products, a change of pressure, adding a catalyst | |
" No Effect | |
"One mole of ethanoic acid reacts with one mole of ethanol at 298K. When equilibrium is reached it is found that two thirds of the acid has reacted. Calculate the value of Kc. CH3COOH(l) + C2H5OH(l) <-------> CH3COOC2H5(l) + H2O(l) | |
" 4 | |
What is the formula for mole fraction? | "number of moles of a substance/number of moles of all substances present |
" | |
What is the formula for partial pressure? | "partial pressure = total pressure x mole fraction |
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What is formula for total pressure of a mixture of gases in terms of its partial pressure. | "total pressure = sum of the partial pressures |
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"A mixture of 16g of O2 and 42g of N2 , exerts a total pressure of 20000 Nm-2. What is the partial pressure of each gas ? | |
" O2 = 5000N/M2 N2 = 15000N/M2 | |
"When nitrogen (1 mole) and hydrogen (3 moles) react at constant temperature at a pressure of 8 x 106 Pa, the equilibrium mixture was found to contain 0.7 moles of ammonia. Calculate Kp N2(g) + 3H2(g) <-------> 2NH3(g) | |
" "1.73 x 10-14 Pa-2 | |
" | |
"1.000 moles of PCl5 vapour are heated to 500 K in a sealed vessel. The equilibrium mixture, at a pressure of 625 kPa, | |
contains 0.600 moles of chlorine. Calculate Kp. PCl5(g) ⇌ PCl3(g) + Cl2(g) | " 352KPa |
For this equilibrium 2 SO2(g) + O2(g) ⇌ 2 SO3(g) calculate Kp if pO2 = 102 kPa, pSO2 = 251 kPa, pSO3 = 508 kPa. | 0.402KPa-1 |
"A mixture of 1.90 moles of H2 and 1.90 moles of I2 were allowed to reach equilibrium at 710 K. The equilibrium mixture | |
was found to contain 3.00 moles of HI. Calculate Kp: H2(g) + I2(g) ⇌ 2 HI(g) | " 56.3 |
Describe CC. | A column is packed with a solid (silica) and a solvent moves down the column |
Describe GC. | A column is packed with a solid or with a solid coated by a liquid, and an inert gas is passed through the column under pressure at high temperature |
Describe TLC. | A plate is coated with a solid (silica on a plastic plate) and a solvent moves up the plate |
In GC, what is the retention time? | The time taken for the substances to be eluted |
What are the general names given to the phases in chromatography? | Stationary and mobile |
What does CC stand for? | Column Chromatography |
What does GC stand for? | Gas Chromatography |
What does TLC stand for? | Thin Layer Chromatography |
What is chromatography? | A process to separate the constituents of a mixture (usually coloured dyes at iGCSE) - separation depends on the balance between solubility in the mobile phase and the retention by the stationary phase |
What is the Rf value and how is it calculated? | Retention factor = (distance moved by the component of a solute)/(distance moved by the solvent) |
Write the reagents and conditions that will convert an alkene to an alkane | H2 and Ni 200oC |
Write the reagents and conditions that will convert an alkene to an alcohol | Conc H3PO4 High T and High P |
Write the reagents and conditions that will convert an alcohol to an alkene | Conc H2SO4 |
Write the reagents and conditions that will convert an alkene to an alkylhydrogensulfate | Conc H2SO4 180oC |
Write the reagents and conditions that will convert an alkylhydrogensulfate to an alcohol | Dilute with water and warm |
Write the reagents and conditions that will convert an alkene to a haloalkane | Halogen and UV light |
Write the reagents and conditions that will convert an alkane to a haloalkane | HX e.g. HBr |
Write the reagents and conditions that will convert a haloalkane to an alkene | KOH dissolved in ethanol, high temperature |
Write the reagents and conditions that will convert a haloalkane to an alcohol | Aqueous NaOH warm |
Write the reagents and conditions that will convert a haloalkane to an amine | NH3 heat under pressure |
Write the reagents and conditions that will convert a haloalkane to a nitrile | KCN dissolved in ethanol and warm |
Write the reagents and conditions that will convert a primary alcohol to an aldehyde or a secondary alcohol to a ketone | K2Cr2O7, H2SO4, heat |
Write the reagents and conditions that will convert an aldehyde to a carboxylic acid | K2Cr2O7, H2SO4, heat |
What reagent do you use to convert an Aldehyde or Ketone to a Hydroxynitrile | HCN |
Write the reagents and conditions that will convert a nitrile to a carboxylic acid | reflux with dilute hydrochloric acid. |
Write the reagents and conditions that will convert a carboxylic acid to an ester | An alcohol, conc H2SO4, and heat |
Write the reagents and conditions that will convert an alcohol to an ester | An carboxylic acid or acyl chloride , conc H2SO4, and heat |
Write the formula of two reagents that will combine a Ketone or an Aldehyde to an alcohol | LiAlH4 or NaBH4 |
To convert an amine to an amide you have to react the amine with a reagent with what functional group? | Acyl Chloride |
Write the reagent that will convert a Nitrile to an Amine | LiAlH4 |
Write the reagents and conditions that will convert benzene to a cyclohexane | H2 Ni catalyst, 200oC |
Write the reagents and conditions that will convert benzene to a nitrobenzene | Conc HNO3 conc H2SO4 50oC |
Write the reagents and conditions that will convert nitrobenzene to a phenylamine | Sn, conc HCl then NaOH |
What catalyst and what functional group are required to convert benzene to a phenyl ketone. | AlCl3 and an Acid Chloride |