Grade 12 Chemistry Exam Notes PDF

Grade 12 Chem Exam Notes Date: Wednesday, January 29, 2025 Time: 8:30-10:30 Worth: 30% Mark Breakdown: Part ​ Type ​ ​ ​ ​ ​ Marks​ Suggested Time A​ Multiple Choice (K)​ ​ ​ 30​ ​ 35 minutes B​ Problem Solvi...

Grade 12 Chem Exam Notes Date: Wednesday, January 29, 2025 Time: 8:30-10:30 Worth: 30% Mark Breakdown: Part ​ Type ​ ​ ​ ​ ​ Marks​ Suggested Time A​ Multiple Choice (K)​ ​ ​ 30​ ​ 35 minutes B​ Problem Solving & Calculations (T) 40​ ​ 50 minutes D​ Short Answer (C)​​ ​ ​ 17​ ​ 20 minutes E Short Answer (A) 13 15 minutes ​ ​ ​ ​ ​ ​ ​ ​ Total​ ​ ​ ​ ​ ​ 100 120 minutes (2 hours) Unit 1: Energy Changes and Rates of Reaction (Ch. 5 and Ch. 6) 5.1 Changes in Matter and Energy - Definitions Kinetic: energy due to motion Potential: stored energy (Ex. Chemical, Nuclear) Thermochemistry: study of energy changes that accompany physical/chem/nuclear changes in matter Calorimetry: experimental technique - measuring energy changes Chemical system: the substance undergoing the change 1.​ Open system: energy and matter can move in/out (ex. Burning a marshmallow) 2.​ Closed system: energy can move in/out 3.​ Isolated system: neither can move in/out Surroundings: the system environment that absorbs or releases energy Endothermic: energy is absorbed by the system -​ Temperature of the surroundings decreases -​ More energy is needed to break bonds than is released by the formation of new bonds. Exothermic: energy is released by the system -​ Temp of surroundings increases -​ More energy is needed to break bonds than is released by formation of new bonds *Example: Ice melting in your hand → endothermic -​ system = ice (physical change) -​ surroundings = hand Formulas + Symbols q = mcΔT -​ Heat (q): amount of energy transferred between substances in joules (J) -​ Mass (m): amount of matter in grams (g) -​ Specific heat capacity (c= 4.18J/g℃): quantity of heat required to raise the temperature 1℃ of a unit of mass in J/g℃ -​ Temperature (T): average kinetic energy of particles in ℃ Δ𝐻𝑠𝑦𝑠𝑡𝑒𝑚= ± ⎸𝑞𝑠𝑢𝑟𝑟𝑜𝑢𝑛𝑑𝑖𝑛𝑔𝑠 ⎸ Enthalpy Change (ΔH): difference in enthalpies (energy absorbed/released to surroundings) of reactants and products during a change -​ For exothermic reaction -​ Law of conservation of energy - energy may be converted from one form to another, or transferred from one set of molecules to another, but the total energy of the system and its surroundings remains the same. -​ Assumptions 1.​ No heat is transferred between calorimeter and the outside environment 2.​ Any heat absorbed by the calorimeter is negligible 3.​ Aqueous solutions have the same density/specific heat capacity as water -​ The enthalpy change in the system is equal to the heat released to the surroundings 5.2 Molar Enthalpies - Definitions Enthalpy/heat (H): total kinetic energy + potential energy of a substance Standard Enthalpy Change: energy change at SATP Molar Enthalpy: energy change associated with physical, chemical, or nuclear change involving ONE MOLE OF A SUBSTANCE Standard Molar Enthalpy: energy change with one mole of a substance SATP Formulas →Calculate molar enthalpy… ΔHx = ΔH/n or ΔH = nΔHx -​ Solution (ΔHsoln) -​ Combustion (ΔHcomb) -​ Vaporization (ΔHvap) -​ Condensation (ΔHcond) -​ Freezing (ΔHfr) -​ Neutralization (ΔHneut) -​ Formation (ΔHf) *Example: -​ ONE mole of hydrogen as it burns has an enthalpy change of -285.8kJ -​ TWO moles of hydrogen as it burns has an enthalpy change -571.6kJ *Since ΔH = nΔHx and ΔH = qwater than nΔHx = mcΔT 5.3 Representing Enthalpy Changes - Definitions Thermochemical equation: an equation indicating the absorption or release of heat Representing Enthalpy Changes: Method #1: Thermochemical equations with energy terms -​ If the reaction is endothermic is required a certain quantity of energy to be supplied to the REACTANTS -​ -​ If the reaction is exothermic, energy is released and listed with the PRODUCTS. -​ Method #2: Thermochemical equations with ΔH values -​ Write a balanced chemical equation and then the ΔH value beside it making sure it is given the correct sign -​ Units in kJ  -​ Method #3: Molar enthalpies of reaction -​ For an exothermic reaction, the standard molar enthalpy is measured by taking into account all energy required to change the system from SATP. -​ Method #4: Potential energy diagrams 5.4 Hess’s Law of Additivity of Reaction Enthalpies - Definitions Hess’s Law: The value of ΔH for any reaction for any reaction that can be written in steps equals the sum of the values of ΔH for each step -​ ΔHtarget = ΣΔHknown Summary -​ Two or more equations with known enthalpy changes can be added together to form a new target equation, then their enthalpy changes may be added together to get the enthalpy change of the target equations -​ If a chemical equation is reversed then the sign ΔH changes -​ If the coefficients of a chemical equation are altered by multiplying or dividing by a constant factor, then the ΔH is altered in the same way 5.5 Standard Enthalpies of Formation - Definitions Standard enthalpy of formation: the quantity of energy associated with the formation of one mole of a substance from its elements in their most stable form and standard states -​ Writing formation equations 1.​ Write one mole of product in the state that has been specified 2.​ Write the reactant elements in their standard states 3.​ Choose equation coefficients for the reactants to give a balanced equation yielding one mole of product ΔH for elements: the standard enthalpy of formation of an element already in its standard state is zero Formulas 1.​ ΔH° = (sum of standard enthalpies of formation of products) - (sum of standard enthalpies of formation of reactants) 2.​ ΔH - nΔHx 6.1 Rate of Reaction - Definitions Chemical kinetics: the area of chemistry that deals with the rate of reactions - ways to make reactions go slower or faster Rate of reaction: the speed at which a chemical change occurs -​ Rate at which product is formed (+ rate) or reactant is consumed ( - rate) -​ generally expressed as a change in concentration per unit of time Average rate of reaction: the speed at which a reaction proceeds over a period of time (often measured as change in concentration of reactant or product over time - mol/L. s) -​ The absolute value of the slope of the secant line drawn between TWO points Theories Measuring Reaction Rates 1.​ Reactions that produce a gas -​ Measured in volume/pressure -​ Faster reaction = the greater the change in volume and pressure 2.​ Reactions that involve ions -​ Form ions = more conductivity -​ Measured as a function of time 3.​ Reactions that change color -​ As a reaction proceeds the intensity of the colored reactant or product may increase or decrease. Formulas + Symbols Average rate of reaction = Change in concentration → r = Δc Change in time Δt Possible units for the rate of reaction → mol or mol L.s L.min R = average rate of reaction Δt = elapsed time Δc = change in concentration 6.3 Rate Laws and Order of Reaction - Definitions Rate law: the rate ‘r’ will always be proportional to the product of the initial concentrations of the reactants, where these concentrations are raised to some exponential values -​ Expressed as r = [X] ^m [Y] ^n -​ m and n are exponential values - show the relationship between rate and initial concentration -​ They do not have to equal the coefficients of a and b in the equations -​ Written as r = k [X] ^m [Y] ^n -​ X and Y are formulas for reactants -​ a X + b Y → products Rate law equation: the relationship among rate, rate constant, initial concentration, and orders of reaction. -​ r = k [X] ^m [Y] ^n -​ K = the rate constant (slope) that's found empirical and valid for only a specific reaction at a specific temp -​ Rate constant: the proportionality constant in the rate law equation Order of reaction: the exponent value that describes the initial concentration of a particular reactant ***RATE DEPENDS ON INITIAL REACTANT CONCENTRATION RAISED TO VARIOUS EXPONENTS*** -​ If the rate d epends on [reactant] ^ 0 it doesn't depend on this reactant  -​ Doubling the initial concentration does not affect the rate (2^0 = 1) -​ If the rate depends on [reactant] ^ 1 -​ Doubling the initial concentration, doubles rate (2^1 = 2) -​ If the rate depends on [reactant] ^ 2 -​ Doubling the initial concentration quadruples rate (2^2 = 4) Overall order of reactant: the sum of the exponent in the rate law equation -​ The order of reaction with respect to NO is 1 -​ The order of reaction with respect to F2 is 1 -​ Overall order = 2 Half-Life of a Reaction: the amount of time required for the concentration of the reactant to decrease to half of its initial concentration -​ One half-life t = 𝑡 1/2 [𝐴]𝑡 = ½ [𝐴]𝑜 -​ Two half-lives 2t = 𝑡 1/2 [𝐴]𝑡 = ½ x ½ [𝐴]𝑜 = ¼ [𝐴]𝑜 -​ Three half-lives 3t = 𝑡 1/2 [𝐴]𝑡 = ½ x ½ x ½ [𝐴]𝑜 = ⅛ [𝐴]𝑜 Formulas + Symbols Integrated Rate Law Equation -​ Expresses the concentration of a reactant as a function of time -​ [𝐴]𝑡 : concentration at some time after the reaction started (mol/L) -​ [𝐴]𝑐 : initial concentration (mol/L) -​ k: rate constant -​ t: time (sec) -​ ln: natural logarithms First Order −𝑘𝑡 ln[𝐴]𝑡 = -kt + ln[𝐴]𝑜 or ln [𝐴]𝑡 /[𝐴]𝑜 = -kt or [𝐴]𝑡 = [𝐴]𝑜 𝑒 -​ The graph will yield a straight line where k = - slope Second Order 1/ [𝐴]𝑡 = kt + 1/ [𝐴]𝑜 -​ The graph will yield a straight line where k = slope Half-Life of First Order -​ 𝑡 1/2 = 0.693/k Half-Life of Second Order -​ 𝑡 1/2 = 1/ [𝐴]𝑜 6.2 Factors Affecting Reaction Rate - Definitions Chemical Nature: -​ The nature of reactants affects the reaction rate -​ Reactions of monatomic ions are extremely fast whereas molecular substances are slower -​ Weaker bonds = faster rate Concentration: -​ If the initial concentration of a reactant is increased the reaction rate increases -​ Concentrated = more chance of collisions = faster and more effective -​ Dilute acids = same reaction at slower rates Temperature: -​ The temperature of the system increases = and the reaction rate increases  -​ Ex. putting cake batter in an oven quickly produces a chemical reaction where the cake rises Presence of a Catalyst: -​ Catalysts accelerate reaction rate -​ Ex. Enzymes act as catalysts in the human body Surface Area: IN HETEROGENEOUS SYSTEMS… -​ The amount of exposed area affects the reaction rate -​ The reaction rate increases proportionally with the surface area -​ Ex. Powdered sugar dissolves quicker than sugar cubes Theories -​ Five Factors Affecting Rate 1.​ Chemical nature of reactants 2.​ Concentration of reactants 3.​ Temperature 4.​ Presence of a catalyst 5.​ Surface area 6.4 Collision Theory and Rate of Reaction - Definitions Activation energy: minimum increase in potential energy of system required for molecules to react Elementary step: a step in a reaction that requires either 1, 2, or 3 particle collisions Reaction mechanism: a series of elementary steps that make up the overall reaction. -​ Ex. car assembly line -​ The worker who controls the rate of production of cars is the slowest -​ Therefore increasing the “concentration” of workers at the slowest step increases the rate -​ Three rules when proposing a mechanism… -​ Each step must be elementary. -​ The slowest step must be consistent with the rate equation -​ The elementary steps must add to the overall equation Rate in determining step: the slowest step in a reaction mechanism Reaction intermediates: molecules that are formed short-lived products in reaction mechanisms. Activated complex: an unstable chemical species containing partially broken and partially formed bonds that possess the maximum potential energy possible. -​ It may reverse back to reactants or continue to for product molecules -​ Exist as peaks on graphs HOW DO SCIENTISTS DETERMINE THE REACTION FOR A MECHANISM? -​ There is a direct correlation between exponents in the rate equation and the equation coefficients in the rate-determining step in the mechanism. -​ Example: the rate equation for 4HBr + O2 → 2 H2O + 2 Br is r = k [HBr] [ O2] -​ The reaction coefficients of the rate-determining step are also 1 for each molecule.s -​ Therefore, 1 HBr + 1 O2 → reaction intermediate (rate determining step) -​ Generally, if the empirically determined rate equation is… -​ r = k [X]^m [Y]^n -​ Then the rate-determining step of the mechanism must be… -​ mY + nY → products or reaction intermediates Theories -​ Molecules are held together with chemical bonds -​ The collision theory is when chemical reactions only occur if energy is provided to break bonds and this source of energy would be kinetic energy -​ Main concepts: -​ A chemical system contains particles that are in constant random motion  -​ The average kinetic energy of a sample is proportional to the temperature of the sample -​ A chemical reaction must involve collisions -​ An effective collision has enough energy and correct orientation so that new bonds are formed -​ Ineffective collisions happen when particles rebound the collision -​ Rate of reaction = frequency of collisions x effective collisions Equation - to FIND rate of reaction RATE = frequency of collision X Fraction of collision that is effective ​ Concentration​ ​ ​ ​ - Nature of reactant ​ Surface area​ ​ ​ ​ - Catalyst ​ Temperature ​ ​ ​ ​ - Temperature GRAPH FOR 2 TEMPERATURES -​ Shows how the distribution of kinetic energies changes when a substance is heated or cooled -​ The higher the temp the more particles there are with higher kinetic energy. 1.​ ENDOTHERMIC GRAPH​ ​ ​ ​ 2. EXOTHERMIC GRAPH GRAPH -​ The molecules have kinetic energy but rarely any potential energy in ENDOTHERMIC. -​ The molecules that have kinetic energy tend to change into potential energy in EXOTHERMIC. -​ More kinetic energy = less potential energy -​ Less kinetic energy = more potential energy  -​ IF molecules with enough kinetic energy approach close enough, they rearrange and form ACTIVATED COMPLEX -​ If the product has higher potential energy then it will have lower kinetic energy in endothermic 6.5 Explaining and Applying Chemical Kinetics - Definitions Threshold energy: the minimum kinetic energy required to convert kinetic energy to activation energy during the formation of the activated complex Heterogeneous catalyst: a catalyst in a reaction in which the reactants and the catalyst are in different physical states Homogeneous catalyst: a catalyst in a reaction in which the reactants and the catalyst are in the same physical state Maxwell Boltzmann Distribution: -​ an effective collision required a minimum energy which gets converted to potential energy - activation energy - as the activated complex is formed -​ Minimum energy = threshold energy -​ distribution of kinetic energy is called Maxwell Boltzmann Distribution -​ vertical axis = number of molecules with particular kinetic energy -​ horizontal axis = different energies Theories The rate at which a reaction occurs depends on two criteria: 1.​ The frequency of collisions 2.​ The fraction of those collisions that are effective (ask if we should know the theoretical effect on concentration, surface area, temperature, and catalysis) Formulas + Symbols Rate Law: r=k[A]^n[B]^m (This describes qualitative rate dependence with respect to concentrations of reactants) → Changes in both the activation energy and temperature have exponential effects on the value of k and the reaction. Unit 2: Structure and Properties (Ch. 3 and Ch. 4) Development of Quantum Theory Timeline 3.1 Early History of Atomic Theories - Quantum Mechanics: study of how light interacts with matter -​ current theory of atomic structure -​ based on wave properties of electrons Bohr Model -​ Rutherford: studied nature of the light that was produced when electrical currents passed through tubes with gaseous elements -​ Passed this light through a spectrometer -​ Light from elements (hydrogen + helium) = dark background + discrete lines -​ Each element had a unique spectrum -​ The wavelength of each line had a specific energy -​ Neils Bohr: suggested electrons could only exist at certain energy levels -​ Explained atomic spectra and periodic table -​ Atoms gained energy = excited, jumped to high levels -​ Atoms lost energy = fall into original levels -​ “Because there were specific differences between energy levels, specific wavelengths of light were seem in the spectrum” -​ Each element had a unique spectrum -​ Light from elements (hydrogen + helium) = dark background + discrete lines -​ Wavelength of each line had a specific energy -​ Periodic Table -​ Each electron orbit of the same shell could only hold so many electrons -​ 1 shell was filled → next level was filled -​ Chemical properties were based on number of electrons in the outer shell Problems? -​ Why are electrons confined to energy levels? -​ Why don't they constantly give off light? -​ As they change direction/accelerate they should give off light -​ Can't explain the spectra of atoms with more than one electrons on the outer shell -​ What was so special about 2 and 8? Electrons as Waves -​ Louis de Broglie: suggested electrons could act like waves + particles - like light -​ Electron beams could be diffracted/bent through a slit-like light -​ Werner Heisenberg: if an electron traveled like a wave there's no way to locate the EXACT position + speed of an electron in the wave - Uncertainty principle -​ Erwin Schrodinger: created wave functions -​ They form regions of space called orbitals - electron-density clouds -​ The densest area of the cloud = greatest probability of finding an electron Wave Functions -​ Three quantum numbers: -​ Principal # (n) - energy level -​ Secondary # (l) - how fast an electron moves in its orbit - the shape of the orbital -​ Magnetic number (𝑚𝑙) - orientation in space -​ Spin (𝑚𝑠)- no two electrons can be in the same state/have the same 4 quantum numbers -​ clockwise/counterclockwise spin of electrons -​ Shapes + numbers -​ S - circle (max orientations = 1) -​ P - dumb-bell (max orientations = 3) -​ D - four lobe (max orientations = 5)  -​ F - six-lobe ( max orientations = 7) -​ Each orbital can only hold 2 electrons -​ Some sublevels begin to overlap in energy - s