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chem unit 2 tgshs
chemistry unit 2
| Term | Definition |
|---|---|
| Surface tension: | The ability of a liquid’s surface to resist external force |
| Effects of surface tension: beading | Beading: Water forms beads on non-polar surfaces like waxy leaves. Water molecules interact more with each other than with the leaf surface. |
| Effects of surface tension: flotation | Flotation: Objects denser than water can float if they don’t interact with water and spread their weight across the surface. This is due to strong hydrogen bonding between water molecules. |
| Effects of surface tension: dripping | Dripping: Water forms droplets on surfaces like taps.As the droplet grows, gravity eventually overcomes hydrogen bonding, causing it to fall. |
| Concentration definition | Measures the amount of solute in a known volume of solution. |
| Concentrated solution definition: | high amount of solute |
| dilute solution definition | low amount of solute |
| saturation definition: | Describes how much solute is dissolved in a solvent |
| Unsaturated solution: | More solute can still dissolve. |
| Saturated solution: | No more solute can dissolve; excess settles at the bottom. |
| Supersaturated solution: | - Formed by heating a saturated solution to dissolve more solute. - If cooled without crystallisation, it remains in solution. - Often used to grow crystals |
| Parts per million (ppm) | - is for measuring very low concentrations - is equivalent to mg/L |
| Solubility Rules | - Not all ionic mixtures form precipitates. - Use solubility rules to predict whether a precipitate will form. s –soluble (>10g/L) p –partially soluble (1-10g/L) i –insoluble (<1g/L) |
| Precipitate reaction | - Occur when two aqueous ionic solutions are mixed. - A precipitate (solid) forms as a result of the reaction |
| Carbonate ions (CO₃²⁻) react with acids to produce... | carbon dioxide (CO₂), water (H₂O), and a salt. |
| what indicated gas production | formation of bubbles |
| testing for carbon | - CO₂ extinguishes flames —test by exposing a lit match to the gas. - Bubbling CO₂ through lime water forms a cloudy suspension. |
| drawing lewis dot structures step 1 | 1.Sum the valence electrons from all atoms shown in the molecular formula. a)For polyatomic anions, add the value of the negative charge to this total. b)For polyatomic cations, subtract the value of the positive charge from this total. |
| drawing lewis dot structures step 2 | 2.Write atomic symbols to show which atoms are connected a)Define central atom.Central atom=most electropositive(or least electronegative)atom(electronegativity table) b)Connect atoms with single line for single pair of bonding electrons |
| drawing lewis dot structures step 3 | 3.Do octets for all outside atoms bonded to central a)Done by adding 3 pairs of electrons to atomic symbol of outside atom (giving each atom total of 8 electrons)(including 2 represented by single bond) Exception is hydrogen(2 electrons in v.s). |
| drawing lewis dot structures step 4 | 4.Place any leftover electrons on the central atom. a)You do this even if it gives the central atom more than an octet. (N.B. There are exceptions to the octet rule.) |
| drawing lewis dot structures step 5 | 5.If there aren’t enough electrons to give the central atom an octet, form multiple bonds. a)You do this by using one or more pairs of non-bonding electrons from the peripheral atoms (choosing the most electropositive, or least electronegative, atom). |
| Exceptions to octet rule | There are three types of ions/molecules that do not follow the octet rule: - Ions or molecules with an odd number of electrons - Ions or molecules with less than an octet - Ions or molecules with more than eight valence electrons(an expanded octet) |
| VSEPR (valence shell electron pair repulsion) theory | Each pair will inhabit space as far away from each other as possible. It predicts shapes of molecules/ions by assuming that valence-shell electron pairs within molecules are arranged around each atom to minimise repulsion between electron pairs. |
| lone pairs in VSEPR | L.p are treated same as b.p cus they take space around central atom=therefore repel other pairs B.p take up space further away from central atom due to being shared with another atom L.p do not experience this force=closer to central atom |
| Linear Shape Properties | - electron configuration - 2 bond pairs - 180° bond angle(s) - Example: BeCl2 |
| Trigonal Planar Shape Properties | - electron configuration - 3 bond pairs - 120° bond angle - Example: AlCl3 |
| Tetrahedral Shape Properties | - electron configuration - 4 bond pairs - 109.5° bond angle - Example: CH4 |
| - electron configuration - 2 bond pairs - 180° bond angle(s) - Example: BeCl2 | Linear Shape |
| - electron configuration - 3 bond pairs - 120° bond angle - Example: AlCl3 | Trigonal Planar Shape |
| - electron configuration - 4 bond pairs - 109.5° bond angle - Example: CH4 | Tetrahedral Shape |
| Irregular Shapes | If a molecule/ion, has lone pairs on central atom, shapes are slightly distorted compared to regular shapes. = extra repulsion caused by lone pairs. As a result of extra repulsion, bond angles tend to be slightly less as the bonds are squeezed together. |
| Molecules with Double Bonds | The shape of a compound with a double bond is calculated in the same way. Double bonds behave exactly as single bonds for repulsion purposes so the shape will be the same as a molecule with two single bonds and no lone pairs. |
| Geometric shape of a molecule | |
| Definition of Solubility | Solubility refers to the maximum amount of solute that can dissolve in a given quantity of solvent at a specific temperature. A solution becomes saturated when no more solute can dissolve. |
| Factors affecting sollubility | - Strength of intermolecular bonds within the substance - Interaction between solvent and solute molecules - Temperature of the solvent |
| Ability to Dissolve in Water | Depends on: - Attraction between solute and polar water molecules - Internal attractive forces within the solute - Temperature and type of solute (gas or solid) |
| Hydrogen Bonding & Energy | - Dissolving solids in water requires disrupting hydrogen bonds between water molecules. - This process needs energy. - Sufficient attraction between solute and water molecules is needed to compensate. |
| Are Molecular Substances Soluble in Water? | - Polar molecular substances can dissolve in water. - Permanent dipoles enable intermolecular forces with water molecules. |
| Non-Polar Molecular Substances | - Interact weakly with water via dispersion forces. - Generally less soluble than polar substances. |
| Polarity and Solubility | - Some solutes have both polar and non-polar regions. - Solubility depends on the size of the non-polar region. |
| Interaction Strength Comparison | Molecular substances: partial charges → weaker interactions Ionic substances: whole charges → stronger interactions |
| Ionic vs. Molecular Solubility | Ionic substances are generally more soluble due to stronger water interactions. |
| Exceptions with Solubility/Insolubility | Some ionic substances are insoluble due to strong internal ionic bonds in their crystal lattice. |
| Temperature and Solubility | - Increasing temperature generally increases the solubility of solids in liquids. - Higher temperature leads to higher kinetic energy of molecules. |
| What is a Solubility Curve? | - A solubility curve is a graph that shows the amount of solute that dissolves in 100 g of water at various temperatures. - Y-axis: Mass of solute - X-axis: Temperature |
| The curve indicates the saturation point at each temperature: | –Below the line: Unsaturated solution (soluble) –On the line: Saturated solution (soluble) –Above the line: Supersaturated solution (insoluble) |
| Solubility of gases | - Increasing temperature raises kinetic energy, breaking solute–solvent bonds. - Atmospheric gases (N₂, O₂, CO₂) have low solubility in water. |
| Temperature Effects on Solubility of gases | - Weak interactions between gas and water break easily. - Dissolved gases evaporate more readily than water. |
| pH definition | - Power of hydrogen - pH is a common measure of ‘acidity’ - It is commonly presented on a 1-14 scale. Where 1 is very acidic and 14 is not acidic/basic/alkaline. |
| Acids definition | - An acid is defined by Arrhenius as a substance which when dissociates in water increases the concentration of hydrogen ions. - The hydrogen ion H+ is also referred to as a proton, or the acidic ion. |
| Bases definition | - An alkaline, base, is defined by Arrhenius as a substance which when dissociates in water increases the concentration of hydroxide ions. - The hydroxide ion OH-is also referred to as the alkaline/ basic ion. |
| pH scale | - a logarithmic scale in relation to the concentration of hydrogen ions. - The higher the hydrogen ion concentration, the lower the number on the pH scale and vice versa. |
| Dissociation | refers to the process of a molecule splitting into smaller parts (ionic substances split apart) |
| pH vs pOH | Both hydrogen and hydroxide ions are present in all samples of water. pH and pOH scales oppose each other. This is important to understand. (eg. [H+]10^0 = [H-]10^-14) |
| how to measure pH | pH =−log[H^+] or [H^+]=10−^pH |
| The behaviour of strong and weak acids and bases in aqueous solutions. | Strong acid/base - increase concentration of hydrogen/hydroxide (respectively) ions (completely dissociate into ions when dissolved in water) Weak acids/bases - partially dissociate |
| which arrows to use with weak and strong acids/bases | strong: one directional arrow (→) representing complete transformation weak: equilibrium arrows (⇌) indication 'dynamic equilibrium', with reaction proceeding both forward and back. No nett change in concentration |
| Strong vs Concentration | just because |
| Acid-base reactions | - produces water and a salt - In neutralisation reaction, hydronium ions (H3O+) combine with hydroxide ions (OH-) to form neutral water. (double displacement) - Neutralisation reactions are one way to produce pure salts. |
| 3 of the strong acids | - hydrochloric acid (HCl), - sulfuric acid (H₂SO₄) - nitric acid (HNO₃) |
| Titration theory | - The concentration of an acid or a base in solution can be determined by performing a neutralisation reaction. - An appropriate acid-base indicator is used to show when neutralisation has occurred. |
| Self ionisation constant of water | Kw = [H^+] x [OH^-] = 10^-14 |
| extra notes on pH and pOH scale | - when at 25°, the product of the concentration of pH and pOH is 10^-14 - sum of logs = 14 |
| What Affects Reaction Rates? | - Temperature - Pressure - Concentration - Surface Area |
| Why Do Reaction Rates Matter? | - Control over chemical processes - Industrial efficiency - Food preservation |
| Collision Theory Basics | - Not all collisions lead to reactions - Reactants must collide: (in order to react) –in the correct orientation –with sufficient energy (activation energy - Eₐ) |
| Activation Energy | - Denoted as Eₐ - Minimum energy required for a successful reaction - Determines reaction success (reaction rate deoens on activation energy size) - Low activation energy = weak bonds, high "" = strong bonds |
| Factors Affecting Activation Energy | - Bond strength (energy) - Number of bonds - Multiple steps for reactions with many bonds |
| What is Maxwell-Boltzmann Distribution? | - Particles in a reaction system have varying energy levels. - Distribution shows number of particles vs. kinetic energy. - X-axis: Kinetic energy (E), Y-axis: Number of particles. |
| Key Features of the Distribution (Maxwell-Boltzmann Distribution) | - Peak of the curve = Modal kinetic energy. - Modal energy is related to the temperature of the reaction. - Area under the curve = Total number of particles. |
| Methods to Increase Reaction Rate | - Adding energy to the system - Manipulating the activation energy - Increasing the frequency of collisions and therefore the proportion of successful collisions |
| 5 things that manipulate reaction rates | ""=affects Concentration: "" liquid, aqueous, gaseous reactants Pressure: "" gaseous reactants only Surface Area: "" solid reactants only Temperature: "" solid, liquid, aqueous, gaseous reactants Catalysts: "" solid, liquid, aqueous, gaseous reactan |
| 1. thing that manipulates reaction rates (concentration) | ^ particles = more frequent collisions b/w reactants + collisions = ^ proportion of successful collisions + successful collisions = faster chemical reaction ^ concentration= ^ reaction rate ↓ concentration = ↓ successful collisions=slow reaction rate |
| 2. thing that manipulates reaction rates (pressure) | - The gas equivalent of concentration - Pressure is force exerted per unit area by one substance on another. - In gases, pressure results from collisions of gas particles with container walls. - ^ number of particles increases the pressure. |
| 3. thing that manipulates reaction rates (surface area) | - Reactions occur at the surface of a solid reactant - ^ surface area = ^ collisions - Finely divided solids have larger surface areas - ^ surface area = the number of successful collisions - More successful collisions = ^ the rate of reaction |
| 4. thing that manipulates reaction rates (temperature) | - Higher temperature → more particles exceed activation energy (E > Ea) - More particles with sufficient energy → more successful collisions - Reaction rate increases with temperature |
| 5. catalyst | |
| Explain how puddles evaporate | - ↑temp =↑KE =↑likelihood of particles equal to Ea =↑proportion of collisions - ↑concentration = particles closer =↑chance of particles colliding - ↑pressure also = ↑chance of particles colliding - Particles @ surface = more likely to evaporate |
| Measuring the Rate of a Chemical Reaction | - Rate = change over time - Observable changes include: –Loss in mass of solid reactant –Volume of gas produced–Formation of precipitate –Colour change –Increase in product concentration –Decrease in reactant concentration |
| rate of reaction unit | The usual unit for the rate of reaction is Ms-1 |
| End point? | - Reaction rate affects the yield of a chemical reaction - Slow reactions may not reach their end point - Monitoring rate is essential for accurate yield prediction |
| Theoretical Yield | - Theoretical yield is calculated using stoichiometry - Requires known amount (mole) of reactants - Used to compare with experimental yield |
| Stoichiometry | - Stoichiometry helps determine reactant-product relationships - Essential for calculating theoretical yield - Based on balanced chemical equations |
| Catalysts | - Catalysts increase the rate of a chemical reaction without being used up. - They provide an alternative pathway with lower activation energy. - Used in the human body, medicines, pharmaceuticals, chemical industry, car exhaust systems, and more. |
| Types of Catalysts | - Homogeneous: Same state as reactants (e.g., aqueous potassium iodide in hydrogen peroxide dissociation). - Heterogeneous: Different state from reactants (e.g., solid iron in ammonia production). |
| How do Catalysts Work? | - Mechanisms of lowering activation energy are key to scientific understanding. -Heterogeneous have high surface energy. –Reactants bond to catalyst surface, weakening internal bonds. –Atoms rearrange on surface to form products. |
| Catalyst Effect on Maxwell-Boltzmann Distributions | - Catalysed reactions shown on Maxwell-Boltzmann distribution with lower activation energy line. - Catalyst lowers activation energy (Ea), increasing particles with energy > Ea. - Proportion of successful collisions increases, but not frequency. |
| Catalyst Effect on Energy Diagrams | - Catalyst provides lower activation energy (Ea) pathway. - Shown as a broken line below uncatalysedpathway in energy profile diagram. - Activation energy (Ec) is lower, but enthalpy change (ΔH) remains unchanged. |
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