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Chem1200 Final Exam Notes.docx

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Department
Chemistry
Course
Chemistry 1027A/B
Professor
Mark Workentin
Semester
Winter

Description
Chapter 1: Chemical Equilibrium 1.1 The Equilibrium Constant  Reversible chemical reactions – at equilibrium, system contains both reactants and products (dynamic)  Equilibrium achieved when rate of forward reaction = rate of reverse  At the molecular level, both forward and reverse processes are still occurring  Equilibrium constant (K) is the mathematical relationship between pressures / concentrations of the reactants and products  Concentration of pure solids and liquids are not included in K  Numerical value of K only changes if the temperature of the system changes  Same phase = homogenous / different phase = heterogeneous  Summary of manipulation rules: √  Magnitude of the equilibrium constant is an indication of the relative amount of product / reactant  If K is large – more products are present, equilibrium lies to the right  If K is small – more reactants are present, equilibrium lies to the left  Reaction Quotient (Q) indicates whether or not the system is at equilibrium and the direction in which the reaction must proceed  Q < K eqreaction must proceed in forward direction  (not enough product) Q > K eqreaction must proceed in reverse direction  (not enough reactant) Q = K : equilibrium eq  Le Châtelier’s Principle: “If chemical system at equilibrium is disturbed, equilibrium will shift in such a way as to minimize the disturbance”  Concentration changes – system will attempt to use additional / make more  Pressure changes – Increase pressure = decrease volume, it will shift to right towards smaller number of moles / decrease pressure = increase volume, it will shift to left towards larger number of moles  If equal number of moles of gaseous reactants and products, position of equilibrium is unaffected by pressure changes  Numerical value of K is not affected by concentration or pressure changes  For exothermic reaction, heat is treated like a product (K decrease) For endothermic reaction, heat is like a reactant (K increase)  ΔG < 0 – reaction will proceed spontaneously  If ΔG is +, equilibrium lies to the left and if ΔG is –, equilibrium lies to the right ( )  Van’t Hoff equation determines K values at different temperatures ( ) 1.2 Solubility of Ionic Compounds  Solubility: amount of substance that will dissolve in a certain volume of a specific solvent  Ionic compounds become more soluble as the temperature increases, completely dissociate into ions once solid dissolves  Solutes dissolve in solvent  Equilibrium is established between undissolved solid and the free ions once no more will dissolve  K sp the equilibrium constant called the solubility product  Excess solid does not affect the position of the equilibrium  Saturated implies that there is an equilibrium between solid and ions, if quantity is less than required it is unsaturated  When Q > K ,spolid will form from ions in solution (supersaturated)  Solution may already contain an ion in common with the dissolving salt – decreases solubility of an ionic compound (K dspsn’t change) 1.3 Weak Acid and Bases  Strong acids and bases completely ionize in solution whereas weak acids and bases do not ionize completely  Arrhenius: Acids produce H O 3ons water while bases produce OH in water  Weak acids and bases proceed until an equilibrium is achieved  Brønsted-Lowry: Acid is proton donor, base is proton acceptor  When a weak acid (HA) ionizes, it donates a hydrogen ion to water  When a weak base (B) ionizes, it accepts a hydrogen ion from water  Lewis Acid: accept a pair of electrons, results in formation of a coordinate covalent bond  Lewis Base: donates a pair of electrons to another atom, forms acid-base adduct  Ammonia and amines behave as bases, metal cations act as acids, and oxides of non-metals behave as acids when reacted with hydroxide  If K > 1, equilibrium lies to the right and K < 1, to the left  Larger the K value, the greater the ionization (stronger = ionize more) +  pK and K share the same relationship as pH and [H ]  If %ionization < 5%, using approximation that c – x ~ c  More dilute the solution of weak acid / base, the greater the percent ionization  Initial concentration is decreased, but in order for Keqo remain constant the % ionization increases (more products relative to reactants at a lower initial concentration) -  A is the conjuage base of the weak acid HA and behaves like a base in solution, HA has one more H and is more positive  BH is the conjugate acid of the weak base B and behaves like an acid in solution  Stronger the weak acid, weaker the weak conjugate base will be  Spectator ions do not affect pH at all  When salt dissolves in water, its cation and anion separate from each other  Amphiprotic anion can act as both acid or base  Strength of weak acid depends on relative thermodynamic stabilities of their conjugate bases  More stable conjugate = stronger the weak will be  Delocalized charges make the species thermodynamically stable and thus, more favorable  Resonance delocalizes charge over two or more atoms and increase stability of the ion (cumulative)  More electronegative atom is better able to bear a negative charge, only works for atoms of same size  Distance effect caused by an electronegative atom stabilizing a negative charge somewhere else on an ion  Larger the area to delocalize, more stable the ion will be  Common ion effect also comes into play with ionization of weak acids and bases  Presence of the common ion = even less weak species will ionize  Amount suppressed by the common ion is pretty significant  Polyprotic: contain more than one ionizable hydrogen atom  K 1> K ;2much more H is produced in the first step than in the second step +  pH of the solution is due almost entirely to the H produced in the first step 1.4 Buffer Solutions  Solution able to withstand changes in pH -  Must contain a weak acid that will react with any OH ions and a weak base that will react with any H 3 ions, and must not react with each other  Amounts of species and conjugate must be roughly equal  Three ways to make an acid buffer: start with both weak acid and conjugate - base, start with only HA and use strong base to convert, start with only A and use strong acid to convert  Analogous combinations can be used to form base buffers  Identifying buffer solutions: identify all species, will there be a reaction, what species are present after  pH of a solution must remain nearly constant, %ionization is fairly small  Value of x is always negligible when compared to the concentration of the parent species  Preferable to use mole amounts  pH of a solution is dependent on the ration of parent to conjugate species  Diluting a buffer solution does not change its pH  Once a buffer solution is formed, it is treated as a equilibrium containing a common ion  Adding a small amount of strong acid / base to a buffer solution results in a small change in pH  Titrations are used to deter+ine amounts -f acid or base in a solution  Equivalence point: moles H = moles OH  Neutralization has occurred but pH is not always 7  Any reaction involving a strong species will go to completion, unidirectional arrow () is used  If one species is weak acid and one is strong, the conjugate of the weak species is produced in the reaction  Acid-base indicators are used to detect equivalence point  Point at which the indicator changes colour is called the endpoint  pK for indicator = pH at equivalence point  Titration curve shows pH resulting from all the reactions occurring  Strong Acid-Strong Base, no buffering action occurs and pH change is very large at the equivalence point (~7)  Weak Acid-Strong Base, magnitude and slope of the increase near the equivalence point is less and contains a buffer (>7)  Weak Base-Strong Acid, equivalence point (<7) and also contains a buffer Chapter 2: Electrochemistry 2.1 Redox Reactions  Oxidation-reduction reactions are referred to as “redox” reactions and occur when there is a change in the oxidation state of one or more elements  Net transfer of electrons from one reactant to another  Consists of two half-reactions; in one electrons are lost (oxidation) and in the other electrons are gained (reduction)  Species that is oxidized is called reducing agent and species that is reduced is called the oxidizing agent  Spectator ions neither gain nor lose electrons  Oxidation and reduction always occur together, number of electrons lost must equal those gained  Oxidation states are used to facilitate the electron accounting in redox reactions (oxidation number) o Pure element = 0 o Equal to the charge of a monatomic ion o Neutral species total must = 0 o Complex ion total must = charge on ion o Fluorine = -1, Group 1 = +1, Group 2 = +2, Hydrogen = +1 except when bonded to a metal, Oxygen = -2, other halogens = -1  When species loses electrons (oxidation), oxidation state increases and when it gains electrons (reduction), oxidation state decreases  Balancing redox reactions o Note changes in oxidation number o Write two half-reactions o Balance coefficients o Add H O to the side deficient in O 2 + o Add number of H to the side deficient in H, for basic solutions add that number of H O2and same number of OH to the other side o Balance charges by adding electrons o Add balance reactions, cancel species  Disproportionation Reaction: converted to both a higher and lower oxidation state as the result of electron transfer  Transition metals can form more than on cation, may take many different oxidation states 2.2 Voltaic Cells  Electrochemistry studies interconversion of chemistry and electrical energy  Electrochemical cell produces an electric current and does electrical work through the use of redox reactions (galvanic cells / voltaic cells)  If a chemical reaction is forced to occur in an electrochemical cell by introducing an electric current from an external source, this is known as electrolysis (electrolytic cell)  Daniell added Zn and Cu in oxidized forms (cations) to the battery in comparable amounts  Overall reaction is carried out by placing a bar of zinc metal in an aqueous solution of copper surface  Electrons will migrate from reducing agent Zn to oxidizing Cu at the solid- liquid boundary, zinc bar will gradually dissolve and metallic copper will appear in the form of a precipitate  Apparatus does not harvest the energy of the electron transfer  Salt bridge is an inverted U-tube filled with aqueous solution of an electrochemically inert salt to balance charge of beaker and limit unnecessary mixing with the reactants  Electrode is where the half-reaction of each half-cell takes place, an electronic conductor in contact with an electrolyte  Electrode where the reduction occurs is called the cathode and electrode where oxidation occurs is called the anode  Cell diagram represents voltaic cell o Anode written first on the left, in the order in which they occur from anode to cathode o Phase boundary is represented by a single vertical bar o Salt bridge is indicated by a double vertical bar o Same phase = separated by comma o Inert electrodes are placed at end and separated by a single vertical bar o Stoichiometric coefficients are not shown  Cell Potential: electrical energy difference between any two electrodes in an electrochemical cell measured using a voltmeter  Voltage measured depends on: nature of reactants, concentration / pressure, and surrounding temperature  Not possible to measure potential for half-cell, only relative to a reference (standard hydrogen electrode = 0 volts)  Standard electrode potential (E°) measures tendency for a reduction process to occur  Stronger oxidizing agents are more easily reduced and have more positive E° redalues (form products that are more difficult to oxidize, product is a weaker reducing agent)  If E°cell0, then reaction is spontaneous as written ( )  Nernst equation describes quantitatively the dependence of a cell potential on the concentrations of reactants and temperature ( ) 2.3 Electrolysis and Electrolytic Cells  Non-spontaneous chemical reaction is forced to occur by application of electrical energy  +ΔG (-E cellll not occur spontaneously  Need to apply enough voltage to overcome the negative sign, to force reaction to go  Direction of electron flow in the cell will be reversed  Solid ionic compound that is not melted / dissolved cannot be electrolyzed because its ions are immobilized and cannot travel  Many active metals are obtained electrolysis of their molten halides  Pure water is difficult to electrolyze, must add an electrolyte (two moles of water react = four moles of electrons)  Overpotential: additional voltage necessary to force a non-spontaneous reaction to be greater than the E°electrolysis  Nature of electrodes and type of chemical reaction determine magnitude of overpotential  For the reduction half-reaction: the higher the value, the more favourable as well as for the oxidation half-reaction  Predicting electrolysis products of mixed molten salts: use ionization energies for cations and use electron affinities for anions  Faraday’s Law: “number of moles of products is equivalent to the number of moles of electrons supplied”  Faraday’s constant is the charge carried by one mole of electron (q e 1.60219 x 10 -19C)  Chlor-Alkali process is the electrolysis of aqueous NaCl that produces sodium hydroxide, hydrogen gas and chlorine gas (applied voltage is only a few volts, an ion-exchange membrane allows Na ions to flow) 2.4 Batteries  Device that converts chemical energy stored in its active materials into electrical energy  Primary batteries can be used only once, irreversible reactions  Secondary batteries can be used, recharged, and reused – easily reversed reactions to regain its charge  Discharging: spontaneous process where electrons migrate from anode to cathode until cell is dead  Zero cell potential = equilibrium  When exhausted cells are recharged, it becomes an electrolytic cell  Lead-Acid battery is the world’s first rechargeable battery and used in today’s cars  Car battery is electrolytic when it recharges, external current is used to provide the energy needed  NiCd has memory effect (after repeated recharging when battery is not completely discharged)  NiMH has no toxic Cd, much less memory effect and store more energy  Li-Ion are very popular due to high energy to weight ratio and do not suffer from memory effect (can rupture de to overcharging)  Li-Pol is a solid phase polymer electrolyte between electrodes and can be stacked in many shapes as possible Chapter 3: Chemical Kinetics 3.1 Reaction Rates and Rate Laws  Study of how quickly a reaction will proceed and factors affecting rate  Speed at which a reaction takes place is governed by several factors: increase concentration, increase temperature, presence of a catalyst increase rate  Rate is defined by comparing the change in the product / reactant concentration over time ( ) ( ) ( ) ( )  Orders tell us how the rate changes when the concentration changes (exponent values)  Products of a reaction do not normally appear in the rate law  Experimental method is used to facilitate monitoring the changes in a reactant or product  First-order reactions is affected linearly  Half-life 1/2) is the amount of time it takes to use up half of the reactant  Concentration or initial amount of the reactant doesn’t affect the half-life ( ) ( )  Zero-order is not dependent on concentration, only on k  Second-order rates depend on k and the square of [A] 3.2 Reaction Mechanisms and the Arrhenius Equation  Reaction coordinate is an abstract one-dimensional coordinate that represents progress along a reaction pathway  Collision theory explains various factors that influence reaction rates  Molecules must overcome the activation barrier, energy required to overcome this barrier is called the activation energy aE )  Average kinetic energy of molecules = temperature  Concentration affects reaction rate (higher concentration = more collisions)  Energy needed to overcome activation barrier comes from heat  Reactants must collide with sufficient energy to overcome the activation barrier, must collide in a proper orientation (steric factor)  Arrhenius equation shows that when the value of E ancreases, the value of k decreases ⁄ ⁄ ⁄ ( ) ( ) ( )  Catalyst is a species that increases the rate of reaction but not consumed in
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