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CHEM_112_Winter Study Guide (Got over 95%)

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Department
Chemistry
Course
CHEM 112
Professor
All Professors
Semester
Fall

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CHEM 112 Winter Study Guide Physical and Chemical Equilibrium Related Formulae For the reaction as follows: , For the reaction in gas phase, , ( ) and thus Definitions is the change in Gibb’s free [ ] is the concentration in mol/L energy P is the partial pressure R is the ideal gas law constant, 8.314 K is the equilibrium constant is the stoichiometric difference between products and reactants T is the temperature in Kelvins Key Concepts Equilibrium A chemical equilibrium displays all of the following characteristics:  The system is closed  The system shows no macroscopic evidence of change  Equilibrium is reached through a spontaneous reaction  A dynamic balance of forward and reverse reactions exists  When temperature is constant, the same equilibrium constant will be obtained, regardless of the direction from which it was approached When K is large, almost 100% of limiting reagent is consumed and thus the reaction is considered to go to completion. Similarly, when K is very small, the reaction is considered not go at all. Types of Equilibrium Systems  Homogenous gas phase systems  Heterogenous mixtures of solid/gas, solid/aqueous, aqueous/gas  Reactions involving substances such as acid/base reactions or redox reactions Relationship of ‘K’ to the Balanced Chemical Reaction  Inverting the BCE causes inversion of K  Multiplying by coefficients by a common factor raises K to that common factor  Dividing coefficients by a common factor causes K to be taken to that factor’s root Le Châtelier’s Principle When an equilibrium system is subjected to a change in temperature, pressure or concentration of reacting species, the system responds by attaining a new equilibrium that partially offsets the impact of the change.  Raising the temperature of the reaction causes the reaction to shift in direction of the endothermic reaction to consume the heat energy  Lowering the temperature of the reaction causes the reaction to shift in the direction of the exothermic reaction to produce the lost energy Example Consider the system as follows: A mixture contains 2.0 x mol of HI, 1.0 x mol of , and 3.0 x mol of in a 2.0L container at 721 K. The equilibrium constant is for this system at 712 K. 1. Is this system at equilibrium? ⁄ ⁄ ⁄ Since , this system is not at equilibrium. 2. If it is not at equilibrium, which way will the system shift to reach equilibrium? Since , the system will shift to the right to reach equilibrium. Electrochemistry Related Formulae , non-standard ( ) , standard conditions Definitions Q is the charge in Coulombs is Faraday’s constant, 96 485 C/mol I is the current in Amperes ( of electrons is potential difference in volts (J/C) Q is the reaction quotient in Nernst eq’n ‘z’ is the # moles of electrons transferred is the electrical work is the change in Gibb’s free energy in the redox reaction Key Concepts Electrochemistry links oxidation-reduction reactions to the production of electrical current. When a chemical loses an electron (LEO), it is oxidized and is termed the reducing agent. When a chemical gains an electron (GER), it is reduced and is termed the oxidizing agent. Disproportionation occurs in redox reactions, when a single substance is both the reducing agent and oxidizing agent. Balancing Redox Reactions Oxidation Number Method 1. Eliminate spectator ions, , and 2. Identify elements which change oxidation numbers 3. Determine increase in oxidation number in reducing agent and decrease in oxidation number in oxidizing agent. 4. Equalize increase in oxidation number with decrease in oxidation number 5. Balance charge by adding 6. Balance hydrogen by adding 7. Replace spectator ions. Half-Cell Method 1. Eliminate spectator ions, , and 2. Divide skeleton equation into two half-cell equations, one for species that are oxidized and the products and one for species that are reduced and the products. 3. Balance all other atoms other than hydrogen and oxygen in each half-cell reaction. 4. Balance oxygen in each half-cell reaction by adding 5. Balance hydrogen in each half-cell reaction by adding 6. Balance charge in each cell by adding electrons to whichever side requires additional negative charge. 7. Balance electrons in each half-cell reaction so that the number of electrons consumed in one half-cell reaction is equivalent to the number of electrons produced in the other half-cell reaction. 8. Add two half-cell reactions to eliminate electrons. 9. Replace spectator ions and simplify. The Galvanic Cell This can be represented schematically by | || | where the double line represents the salt-bridge, that enables net flow of positive ions through it in the right-hand beaker, and of negative ions into the left-hand beaker to preserve charge neutrality in each solution. If the electrochemical reaction occurs spontaneously, it is a galvanic cell. If the reaction operates opposite to spontaneous direction, it is an electrolytic cell capable of recharging. Standard Cell Potentials The standard cell potential measures the tendency for reduction to occur at an electrode. To determine cell potential for an electrochemical cell: 1. Find two half-cell reaction making up a redox reaction and find E value for half reactions from table of standard reduction potentials. 2. Reverse cell reaction with lower reduction potential. This will be the oxidation half-reaction. 3. Balance electrons in two half-cell reactions. (note: E values will remain unchanged) 4. Find sum of two half reactions and value for the cell. 5. For a redox reaction to be spontaneous, must be positive. Acids and Bases Related Formulae For the reaction , pH + pOH = 14 Defintions : ion product of water, 1.0 x : acid ionization constant : base ionization constant Key Concepts Arrhenius suggested that an acid is a substance that when dissolved in water increases hydrogen ion concentration. A base by this definition would increase hydroxide concentration when dissolved in water. Bronsted-Lowry proposed that acids are proton (hydrogen ion) donors and bases are proton (hydrogen ion) acceptors. Acids are bases by this definition exist as conjugate pairs that differ only in one hydrogen atom.  Strong acids: HCl, HBr, HI, , , , o Conjugates of these compounds are weak bases and will not dissociate fully in water, but rather to the point of equilibrium determined by the base ionization constant  Strong bases: NaOH, KOH, RbOH, CsOH, , , o Conjugates of these compounds are weak acids and will not dissociate fully in water, but rather to the point of equilibrium determined by the acid ionization constant Hydrolysis  Salts of strong acids and strong bases do not hydrolyze  Salts of weak acids and strong bases hydrolyze in basic conditions  Salts of weak bases and strong acids hydrolyze in acidic solutions  Salts of weak bases and weak acids hydrolyze o Cations are acidic, anions are basic, solution acidity/basicity depends on relative values of Ka and Kb Common Ion Effect Buffer Solutions A buffer solution is any solution that resists changes in pH even with small additions in acid or base. A buffer solution is made by adding weak acid to a salt containing its conjugate base or by adding a weak base to a salt containing its conjugate base. Optimal buffers have the concentrations of the weak acid and conjugate base (or vice versa) as equal as possible and contain an acid with a value close to the desired pH. Titration Endpoint occurs when the acid-base indicator changes colour indicating that the stoichiometric amount of first reactant has been added to the second reactant. The equivalence point is the point at which the concentration of titrant is stoichiometrically equivalent to that of the unknown solution. At the equivalence point, the pH is not necessarily 7, as the salt of produced can be acidic or basic. In this titration curve, a weak acid is being titrated with a strong base. At each point, the pH can be calculated using the Henderson Hasselbach equation with the knowledge of what concentrations are present. Solution Equilibrium Related Formulae For the reaction, , Note: the AgI is not included in the equilibrium expression, as it is in solid phase. Definitions is the solubility product Key Concepts As with previous equilibrium concepts, precipitation will depend on the reaction quotient, Q, in relation to the solubility product constant, . Ion Solubility Exceptions Soluble None Soluble None Soluble , , Soluble , , Insoluble , , , , Insoluble Group IA, Insoluble Group IA, Insoluble Group IA, , , Soluble None Soluble None Soluble None
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