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Equilibrium = the point in a chemical reaction where the reactants and the products are formed and broken at the same rate
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Dynamic equilibrium = a balance between the forward and backward rates that are occurring simultaneously
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Equilibrium law = mathematical description of a chemical system at equilibrium
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Equilibrium constant (K or K
eq) = the numerical value defining the equilibrium law for a system at a given temperature (changes with temperature) |
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Heterogeneous equilibrium = products and reactants are in at least 2 different states; pure solids/liquids are not included in K
eq formula |
eq)
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K
eq formula: |
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If a[A] + b[B] ⇌ c[C] + d[D]; then
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K
eq = ([C]c[D]d) ÷ ([A]a[B]b) |
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Magnitude of K
eq: states whether the equilibrium position favours products/reactants |
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If K = 1
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[products] = [reactants]
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If K1
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[products] [reactants]
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If K1
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[products] [reactants]
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K
forward vs Kbackward |
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If a[A] ⇌ b[B], then
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K
forward = ([B]b) ÷ ([A]a) |
K
backward = ([A]a) ÷ ([B]b) |
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So ∴ K
forward = 1/Kbackward @ equilibrium |
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Purpose of K
eq: to determine equilibrium concentration of chemical entities given initial conditions (I.C.E. table) |
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Helps to determine the position of the equilibrium of a system using the rate law for the system and comparing it with the K
eq |
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If Q < K
eq |
[products] < [reactants]
Reaction has not reached ⇌ yet; reaction needs to shift right |
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If Q > K
eq |
[products] > [reactants]
Reaction has not reached ⇌ yet; reaction needs to shift left |
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If Q = K
eq |
[products] = [reactants]
Reaction has not reached equilibrium yet; no shift will occur |
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Le Châtelier's Principle: When a chemical system at equilibrium is disturbed by a change in property, the system responds in a way that opposes the change
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Concentration/Temperature
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[conc]/T = shift to consume
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[conc]/T = shift to replace
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If you add more reactant/heat to a system, the system will consume it to make more product, and vice versa
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If you remove reactant/heat from a system, the system will replace it from the existing product, and vice versa
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Volume/Pressure (gases only)
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V =P = shift toward side with more gas entities (i.e. more mol of gas) (more space for particles)
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V =P = shift toward side with fewer gas entities (i.e. less mol of gas) (less space for particles)
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Catalysts/Inert (noble) gases No effect
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Molar solubility: The amount of solute (in mol) that can be dissolved in 1 L of solvent at a certain temperature
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T α solubility: as temperature increases, so does molar solubility
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All compounds have some solubility, even if they are considered "insoluble" (really very low molar solubility)
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Very soluble compounds (high molar solubility) = no ⇌ (complete disassociation into ions)
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Slightly soluble compounds (low molar solubility) = has ⇌ (incomplete/partial disassociation into ions)
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As a compound is placed in a solvent, some part of it will dissolve, but at the same time, the reverse reaction starts
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Eventually the two reactions reach equilibrium, creating a saturated solution (at this point, [conc] of ions remains constant) solubility equilibrium
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Equilibrium formula for solubility:
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If x[A
aBb] (s) ⇌ a[Ab+] (aq) + b[Ba-] (aq), then |
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K
sp = [Ab+]a[Ba-]b |
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K
sp (Solubility product constant): the product of the [conc] of ions in a saturated solution |
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If K
sp1 |
Amount (mol) of aqueous ionsamount (mol) of solid substance
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If K
sp1 |
Amount (mol) of aqueous ionsamount (mol) of solid substance
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Molar solubility and the K
sp describe the solubility of a substance in different ways, meaning you can use one to solve for the other (using an ICE table) |
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Using Q and K
sp to predict precipitate formation |
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If Q < K
sp |
Not enough ions in solution (unsaturated); no precipitate; reaction will shift to the right
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If Q > K
sp |
Lots of ions in solution (saturated); precipitate can form; reaction will shift to the left
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If Q = K
sp |
System is at ⇌; no precipitate (saturated solution)
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pH/pOH: The measure of acidity/alkalinity of a solution
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pH: The measure of [H+] in a solution
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pH = -log[H+]
= -log[H 3O+] |
[H
3O+] = [H+]= 10-pH |
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pOH: The measure of [OH¯] in a solution
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pH = -log[OH¯]
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[OH¯] = 10-pOH
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The pH and pOH values are related to the exponent of K
w (14): |
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pH + pOH = 14
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Arrhenius acid/base = solution that ionizes into H+ (acid)/OH¯ (base) ions
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Bronsted-Lowry acid/base = solution that donates (acid)/receives (base) H+ ions
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Strong acid/base: solution that completely ionizes (acid)/disassociates (base) into ions
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Weak acid/base: solution that partially ionizes (acid)/disassociates (base) into ions
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Monoprotic acids: acids that donate one H+ ion
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Polyprotic acids: acids that donate more than one H+ ion
(diprotic = 2 H+, triprotic = 3 H+, etc.) |
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Amphiprotic substance: substance that can behave like an acid or as a base (i.e. can donate and receive H+ ions)
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Neutralization reactions
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With strong acids/bases:
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acid + base salt + water
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*Complete ionization, so no equilibrium analysis*
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With weak acids/bases:
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acid + base ⇌ conjugate base + conjugate acid
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*Partial ionization, so equilibrium has to be analyzed*
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Acid/Base Constants (K
a and Kb) |
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K
a/Kb indicate the strength of an acid/base |
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If K
a/Kb1 |
Strong acid/base (complete ionization/disassociation)
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If K
a/Kb1 |
Weak acid/base (partial ionization/disassociation)
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If one K is known, the other can be determined using K
w through the formula: |
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K
w = Ka (of acid) ⋅ Kb (of conjugate base) |
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K
w = Kb (of base) ⋅ Ka (of conjugate acid) |
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w)
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Water can dissociate into ions on its own:
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H
2O(l) ⇌ H+(aq) + OH¯(aq) |
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But the H+ ion can also attack other H
2O molecules: |
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H
2O(l) + H+(aq) ⇌ H3O+(aq)1 |
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Adding both reactions together:
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2 H
2O(l) ⇌ H3O+(aq) + OH¯(aq) |
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All equilibria have a constant (K) value, therefore:
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K
w = [H3O+][OH¯](H 2O(l) not included because it is not (aq)) |
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Since water is neutral (pH = 7):
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[H+] = 1.0⋅10-7 [H
3O+] = 1.0⋅10-7 |
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pH + pOH = 14 pOH = 7
[OH¯] = 1.0⋅10-7 |
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If [H
3O+] = 1.0⋅10-7, and [OH¯] = 1.0⋅10-7, then: |
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K
w = (1.0⋅10-7)(1.0⋅10-7)= 1.0⋅10-14* |
3O+(aq): Hydronium ionw is always 1.0⋅10-14
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Cheatography n. The art of mapping complex subjects into clear, instantly navigable reference guides.