How do you write the Kc expression for aA + bB ⇌ cC + dD?
Answer: Kc = [C]^c[D]^d / ([A]^a[B]^b)
Explanation
Pure solids/liquids (and solvent) are omitted. Example: N2+3H2⇌2NH3 → Kc=[NH3]^2/([N2][H2]^3)
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AP Chemistry: Equilibrium, Acids and Bases (High School)
Equilibrium ExpressionsQ vs KICE TablesLe Chatelier's PrincipleKsp and SolubilityStrong Acids & BasesKa, Kb, Kw RelationspH and pOHPercent IonizationBuffersTitration CurvesSalt Hydrolysis
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Equilibrium Expressions
How do you write the Kp expression for a gas-phase reaction?
Answer: Kp = (P_products)^coeff / (P_reactants)^coeff, using partial pressures of gases only.
Explanation
Example: for N2+3H2⇌2NH3, Kp = P_NH3^2 / (P_N2 · P_H2^3)
How do Kc and Kp relate?
Answer: Kp = Kc(RT)^Δn, where Δn = (moles gas products) − (moles gas reactants), R=0.0821 L·atm/(mol·K)
Explanation
For N2+3H2⇌2NH3, Δn=2−4=−2, so Kp=Kc(RT)^−2
Q vs K
What is the reaction quotient Q?
Answer: Q has the same mathematical form as K but is calculated using concentrations/pressures at any point, not necessarily at equilibrium.
Explanation
Comparing Q to K tells you which way the system must shift to reach equilibrium.
If Q < K, which way does the reaction shift? If Q > K?
Answer: Q<K → shifts forward (toward products); Q>K → shifts reverse (toward reactants); Q=K → already at equilibrium.
Explanation
Example: Q=0.5, K=2 → Q<K, so more product forms.
ICE Tables
What do the rows of an ICE table represent?
Answer: I = Initial concentrations, C = Change (±x × stoichiometric coefficient), E = Equilibrium (I+C); substitute E row into the K expression to solve for x.
Explanation
Example: A⇌B, start [A]0=1.0 M, [B]0=0 → change −x/+x → equilibrium (1.0−x), x.
Le Chatelier's Principle
Le Chatelier: effect of adding more reactant?
Answer: Equilibrium shifts right (toward products) to partially consume the added reactant.
Explanation
Example: adding N2 to N2+3H2⇌2NH3 pushes the reaction toward more NH3.
Le Chatelier: effect of removing product?
Answer: Equilibrium shifts right (toward products) to replace the product that was removed.
Explanation
Removing NH3 as it forms drives the reaction further forward.
Le Chatelier: effect of decreasing volume (increasing pressure)?
Answer: Equilibrium shifts toward the side with fewer total moles of gas.
Explanation
For N2+3H2⇌2NH3 (4 mol gas → 2 mol gas), compressing the system shifts it toward NH3.
Le Chatelier: effect of adding an inert gas at constant volume?
Answer: No shift — partial pressures and concentrations of reactants/products are unchanged even though total pressure rises.
Explanation
Only total pressure increases; each species' own partial pressure stays the same.
Le Chatelier: effect of raising temperature?
Answer: Shifts in the endothermic direction, and the value of K itself changes (heat acts like a reactant or product).
Explanation
For an exothermic forward reaction, heating shifts the equilibrium left and decreases K.
Le Chatelier: effect of a catalyst on equilibrium position?
Answer: No shift in position — a catalyst speeds up forward and reverse rates equally, only shortening time to reach equilibrium.
Explanation
K and equilibrium concentrations are unaffected by a catalyst.
Ksp and Solubility
Define Ksp.
Answer: The equilibrium constant for dissolution of a slightly soluble ionic solid: for MmXn(s)⇌mM^n+(aq)+nX^m−(aq), Ksp=[M^n+]^m[X^m−]^n
Explanation
Example: AgCl(s)⇌Ag++Cl−, so Ksp=[Ag+][Cl−]
How do you find molar solubility (s) from Ksp for a salt like CaF2 (AB2 type)?
Answer: For AB2⇌A^2++2B^−, Ksp=(s)(2s)^2=4s^3; solve for s=(Ksp/4)^(1/3)
Explanation
Example: if Ksp=4×10^−11, then s=(1×10^−11)^(1/3)≈2.2×10^−4 M
What is the common ion effect?
Answer: Adding a soluble salt that shares an ion with a slightly soluble solid shifts its dissolution equilibrium left, decreasing solubility.
Explanation
Example: AgCl is less soluble in 0.1 M NaCl solution than in pure water because of the added Cl−.
Strong Acids & Bases
List the 6 common strong acids to memorize.
Answer: HCl, HBr, HI, HNO3, H2SO4, HClO4 (HClO3 is also often included).
Explanation
These ionize essentially 100% in water, so [H+] ≈ initial acid concentration.
Which bases are considered strong (memorize)?
Answer: Group 1 hydroxides (LiOH, NaOH, KOH, RbOH, CsOH) and heavier Group 2 hydroxides (Ca(OH)2, Sr(OH)2, Ba(OH)2).
Explanation
They dissociate completely in water; Ca(OH)2 etc. give 2 mol OH− per mol dissolved.
Ka, Kb, Kw Relations
Write the Ka expression for a weak acid HA.
Answer: For HA⇌H++A−, Ka=[H+][A−]/[HA]
Explanation
Smaller Ka means a weaker acid (less ionization).
Write the Kb expression for a weak base B.
Answer: For B+H2O⇌BH++OH−, Kb=[BH+][OH−]/[B]
Explanation
Smaller Kb means a weaker base.
What is Kw and its value at 25°C?
Answer: Kw=[H+][OH−]=1.0×10^−14 at 25°C, from the autoionization of water.
Explanation
Use Kw to convert between [H+] and [OH−]; e.g., [H+]=1×10^−3 → [OH−]=1×10^−11.
How are Ka and Kb related for a conjugate acid-base pair, and what is pKa?
Answer: Ka×Kb=Kw (=1.0×10^−14 at 25°C); pKa=−log(Ka), with smaller pKa meaning a stronger acid.
Explanation
Example: Ka(NH4+)=5.6×10^−10, Kb(NH3)=1.8×10^−5, product ≈1.0×10^−14; pKa(NH4+)≈9.25
pH and pOH
How do you calculate pH?
Answer: pH = −log[H+] (or [H3O+])
Explanation
Example: [H+]=1×10^−3 M → pH=3
How do you calculate pOH?
Answer: pOH = −log[OH−]
Explanation
Example: [OH−]=1×10^−5 M → pOH=5
What is the relationship between pH and pOH at 25°C?
Answer: pH + pOH = 14
Explanation
Example: pH=3 → pOH=11
How do pH values classify a solution as acidic, basic, or neutral at 25°C?
Answer: pH<7 → acidic; pH=7 → neutral; pH>7 → basic.
Explanation
Based on [H+] vs [OH−] relative to 1×10^−7 M each at 25°C.
Percent Ionization
How do you calculate percent ionization of a weak acid, and how does it change with dilution?
Answer: % ionization = ([H+]ionized / [HA]initial) × 100%. It increases as the initial concentration decreases (more dilute → larger fraction ionizes).
Explanation
Example: if 0.01 M of a 1.0 M acid ionizes, % ionization = 1%; diluting the acid raises this percentage even though [H+] itself drops.
Buffers
What is a buffer solution?
Answer: A solution that resists pH changes, typically a weak acid plus its conjugate base (or weak base plus conjugate acid) present in comparable amounts.
Explanation
Example: CH3COOH/CH3COO− buffer resists added H+ or OH−.
State the Henderson-Hasselbalch equation and what determines buffer capacity.
Answer: pH = pKa + log([A−]/[HA]). Buffer capacity depends on the absolute concentrations of HA and A− — higher, more equal amounts resist larger pH changes.
Explanation
Example: pKa=4.74, [A−]=[HA] → pH=4.74 (most effective buffering occurs near pH≈pKa).
Titration Curves
What defines the equivalence point in a titration?
Answer: The point where moles of acid added equal moles of base (complete stoichiometric neutralization), marked by the steepest part of the curve.
Explanation
For a strong acid-strong base titration, the equivalence point pH = 7.
What is special about the half-equivalence point in a weak acid titration?
Answer: At half-equivalence, [HA]=[A−], so by Henderson-Hasselbalch, pH = pKa.
Explanation
This lets you read a weak acid's Ka directly off its titration curve.
Identify these two titration curves: which is strong acid–strong base and which is weak acid–strong base?
Answer: The curve with equivalence point pH=7 and a very sharp jump is strong acid–strong base; the curve with a gradual buffer region, equivalence point pH>7, and smaller jump is weak acid–strong base.
Explanation
Example: titrating HCl with NaOH gives equivalence pH=7; titrating CH3COOH with NaOH gives equivalence pH≈8.5.
How do you choose an appropriate indicator for a titration?
Answer: Pick an indicator whose color-change range (≈ pKa ± 1) includes the pH at the equivalence point.
Explanation
Example: phenolphthalein (range ~8.2–10) suits a weak acid–strong base titration with equivalence pH>7; methyl orange suits strong acid titrations near pH 7 or lower.
Salt Hydrolysis
Is a salt from a strong acid and strong base acidic, basic, or neutral?
Answer: Neutral — neither ion hydrolyzes water.
Explanation
Example: NaCl dissolves to give Na+ and Cl−, both spectator ions.
Is a salt from a weak acid and strong base acidic, basic, or neutral?
Answer: Basic — the conjugate base of the weak acid hydrolyzes water to produce OH−.
Explanation
Example: NaCH3COO solution is basic: CH3COO− + H2O ⇌ CH3COOH + OH−
Is a salt from a strong acid and weak base acidic, basic, or neutral?
Answer: Acidic — the conjugate acid of the weak base hydrolyzes to produce H+.
Explanation
Example: NH4Cl solution is acidic: NH4+ ⇌ NH3 + H+
Is a salt from a weak acid and weak base acidic, basic, or neutral?
Answer: Depends on relative Ka and Kb of the ions: Ka>Kb → acidic; Kb>Ka → basic; Ka=Kb → neutral.
Explanation
Example: NH4CH3COO is nearly neutral because Ka(NH4+) ≈ Kb(CH3COO−).
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