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Energetics
AQA A Level Chemistry
| Question | Answer |
|---|---|
| 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 |
| Write an equation for the reaction that represents the enthalpy change of combustion of ethanol (C2H5OH). | C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(l) |
| 1.24 g of ethanol was used to heat 50.0 g of water in copper calorimeter The temperature of the water rose by 38°C. Calculate the enthalpy of combustion of ethanol determined by this experiment. The specific heat capacity of the solution is 4.18 J K–1 g–1 | –295 kJ mol–1 |
| Suggest three reasons why the value obtained in an experiment to find the enthalpy of combustion of ethanol, by heating a copper calorimeter is less exothermic than a data book value | 1. heat loss to the air and heating up the equipment 2 incomplete combustion 3 some ethanol evaporate |
| 1.15 g of zinc reacted with 50.0 cm3 of 0.500M CuSO4. The temperature rose by 17.0°C. Which is the limiting reagent and then find the enthalpy change for this reaction. Density is 1.00 g cm–3 c is 4.18 J K–1 g–1 CuSO4(aq) + Zn(s) → Cu(s) + ZnSO4(aq) | –202 kJ mol–1 |
| Calculate the standard enthalpy for the reaction C2H4(g) + H2(g) --> C2H6(g) Given the following standard enthalpy changes of combustion ΔHc C2H4(g)= -1411KJmol-1 ΔHc H2(g)= -286KJmol-1 ΔHc C2H6(g) = -1560KJmol-1 | -137KJmol-1 |
| Calculate the standard enthalpy change of combustion of ethene C2H4(g) Given the following standard enthalpy changes of formation ΔHf C2H4(g)= -52KJmol-1 ΔHf CO2(g)= -394KJmol-1 ΔHf CO2(g) )= -286KJmol-1 | -1412KJmol-1 |
| Calculate the standard enthalpy change of formation of methane CH4(g) Given the following standard enthalpy changes of combustion ΔHcCH4(g)= -890KJmol-1 ΔHc H2(g)= -286KJmol-1 ΔHcC2(s) )= -394KJmol-1 | -76KJMol-1 |
| Calculate the standard enthalpy for the reaction CaCO3(s) --> CaO(s) + CO2(g) Given the following standard enthalpy changes of formation ΔHf CaCO3(s)= -1207KJmol-1 ΔHf CaO(s)= 635KJmol-1 ΔHf CO2(g) = -394KJmol-1 | 178KJmol-1 |
| Calculate the standard enthalpy of formation of solid ammonium chloride from the following data NH4Cl(s) NH3(g) + HCl à NH4Cl(s) ΔHr = -176KJmol-1 ΔHf NH3(g)= -46.1KJmol-1 ΔHf HCL(g)=92.3KJmol-1 | -314.4KJmol-1 |
| Calculate the enthalpy of combustion of propane, C3H8(g), given the following data. ΔfH C3H8(g) = –104 kJ mol–1 ΔcH C(s) = –394 kJ mol–1 ΔcH H2(g) = –286 kJ mol–1 | ΔcH = –2222 kJ mol–1 |
| Pentane burns well in oxygen. C5H12(l) + 8O2(g) → 5CO2(g) + 6H2O(l) Calculate the enthalpy change for this reaction given the following enthalpies of formation: ΔfH / kJ mol–1 C5H12(l) = –147 CO2(g) = –394 H2O(l) = –286 | –3539 kJ mol–1 |
| 1.56 g of pentane was burned and used to heat 100.0 g of water The temperature of the water rose by 28°C. Calculate the enthalpy of combustion of pentane determined by this experiment. The specific heat capacity of the solution is 4.18 J K–1 g–1 . | –540 kJ mol–1 |
| Calculate the enthalpy change for this reaction using the enthalpies of formation shown. TiCl4(s) + 2Mg(s) → 2MgCl2(s) + Ti(s) ΔfH / kJ mol–1 TiCl4(s) = –912 MgCl2(s) = –642 | ΔH = –372 kJ mol–1 |
| Ethene reacts with hydrogen as shown: CH2=CH2(g) + H2(g) → CH3CH3(g) ΔH = –99 kJ mol–1 Calculate the bond enthalpy for the C=C bond using this data and the following bond enthalpies. C-H = 413, H-H = 463, C-C = 348 kJ mol–1 | C=C = 612 kJ mol–1 |
| 1.22 g of propan-1-ol, C3H7OH(l), was burned and used to heat 50.0 g of water The temperature of the water rose by 52°C. Calculate the enthalpy of combustion of propan-1- ol The specific heat capacity of the solution is 4.18 J K–1 g–1 . | –535 kJ mol–1 |
| Hydrogen reacts with oxygen as shown: H2(g) + ½O2(g) → H2O(l) ΔH = –205 kJ mol–1 Calculate the bond enthalpy for the O=O bond using this and this data. Bond enthalpies: H-H = 463, O-H = 436 kJ mol–1 Enthalpy of vaporisation of water = +44 kJ mol–1 | +496 kJ mol–1 |
| Calculate the enthalpy of formation of propan-1-ol, C3H7OH(l), given the following data. ΔcHC3H7OH(l) = –2010 kJ mol–1 ΔcH C(s) = –394 kJ mol–1 ΔcH H2(g) = –286 kJ mol–1 | ΔfH = –316 kJ mol–1 |
| 1.08 g of methanol, CH3OH(l), was burned and used to heat 100.0 g of water The temperature of the water rose by 38°C. Calculate the enthalpy of combustion of methanol . The specific heat capacity of the solution is 4.18 J K–1 g–1 | –470 kJ mol–1 |
| Define Standard enthalpy change of atomization | the enthalpy change when 1 mol of atoms in the gas phase is formed from its element in its defined physical state under standard conditions (298.15K, 1 atm). |
| Define Standard enthalpy change of solution | the enthalpy change observed when one mole of an solute is dissolved completely in an excess of solvent under standard conditions. |
| Define Standard enthalpy change of fusion | the enthalpy change required to completely change the state of one mole of substance between solid and liquid states under standard conditions. |
| Define Standard enthalpy change of vapourisation | the enthalpy change required to completely change the state of one mole of substance between liquid and gaseous states under standard conditions. |
| Define Lattice Enthalpy | the enthalpy required to separate one mole of an ionic compound into separated gaseous ions to an infinite distance apart (meaning no force of attraction) under standard conditions. |
| Define the first electron affinity | The first electron affinity is the energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous -1 ions. |
| Define The first ionization energy | The 1st ionization energy of the element M is a measure of the energy required to remove one electron from one mole of the gaseous atoms M to form one mole of gaseous ions each with a +1 charge |
| Define The average bond dissociation enthalpy | The average bond dissociation enthalpy is the energy needed to break one mole of the bond to give separated atoms - everything being in the gas state. |
Created by:
davenantchemistry
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