Chemical Equilibrium Flashcards

Flashcards covering key vocabulary and concepts related to chemical equilibrium.

Brønsted-Lowry base

A species that accepts an H+.

Base-dissociation constant (Kb)

The equilibrium constant for the reaction of a weak base with water.

pKb

− log 𝐾b, used as an indication of base strength (stronger base ⟹ smaller pKb).

Anions of Weak Acids

Often function as weak bases, accepting a proton from water to produce hydroxide ions.

Relationship between Ka and Kb for a Conjugate Acid-Base Pair

𝐾a × 𝐾b = Kw = 1.0 × 10−14 and pKa + pKb = 14.00

Salts that Yield Neutral Solutions

Consist of the cation of a strong base and the anion of a strong acid. Neither ion reacts with water to any great extent.

Salts that Yield Acidic Solutions

Consist of the cation of a weak base and the anion of a strong acid. The cation reacts with water to produce H3O+.

Salts that Yield Acidic Solutions (Metal Cations)

Consist of a small, highly charged metal cation and the anion of a strong acid. The hydrated metal ion reacts with water to produce H3O+.

Salts that Yield Basic Solutions

Consist of the anion of a weak acid and the cation of a strong base. The anion reacts with water to produce OH–.

Salts of Weakly Acidic Cations and Weakly Basic Anions

The pH of the solution depends on the relative acid strength (Ka) or base strength (Kb) of the ions. If Kb > Ka, the solution will be basic.

The Leveling Effect

All strong acids and bases are equally strong in water because they dissociate completely to form H3O+ or OH–.

Lewis Base

Any species that donates an electron pair to form a bond.

Lewis Acid

Any species that accepts an electron pair to form a bond.

Weak Acid

An acid that dissociates slightly to form ions in water, with most HA molecules remaining undissociated.

Acid Dissociation Constant (Ka)

𝐾a = [H3O+][A−]/[HA], indicating the strength of an acid. Stronger acid ⟹ larger 𝐾a ⟹ smaller p𝐾a.

pKa

-log(Ka)

Types of Weak Acids

HF, acids where H isn't bonded to O or a halogen (HCN, H2S), oxoacids (HClO, HNO2, H3PO4), and organic carboxylic acids (RCOOH).

Solving Weak-Acid Equilibria

1. Write a balanced equation
2. Write an expression for Ka.
3. Define x as the change in concentration.
4. Construct a reaction table in terms of x.
5. Make assumptions that simplify the calculation.
6. Substitute values into the Ka expression and solve for x.
7. Check that the assumptions are justified.

The notation system

• Molar concentrations are indicated by [ ].
• A bracketed formula with no subscript indicates an equilibrium concentration.

The assumptions

• [H3O+] from the autoionization of H2O is negligible.
• A weak acid has a small Ka and its dissociation is negligible. [HA] ≈ [HA]init.

Polyprotic Acid

An acid with more than one ionizable proton. Each dissociation step has a different Ka value, with Ka1 ≫ Ka2 ≫ Ka3.

Acid Strength of Oxoacids

Depends on the electronegativity of the central nonmetal (E) and the number of O atoms around E. Strength increases with electronegativity and number of O atoms.

Arrhenius Acid

A substance with H in its formula that dissociates to yield H+.

Arrhenius Base

A substance with OH in its formula that dissociates to yield OH– .

Brønsted-Lowry Acid

A proton donor, any species that donates an H+ ion. Must contain H in its formula.

Brønsted-Lowry Base

A proton acceptor, any species that accepts an H+ ion. Must contain a lone pair of electrons to bond to H+.

Conjugate Acid-Base Pair

Two species that differ by a proton (H+). The conjugate acid has one more H and one fewer negative charge than its conjugate base.

Net Direction of Reaction

An acid-base reaction favors the formation of the weaker acid and weaker base. The stronger the acid, the weaker its conjugate base.

Autoionization of Water

Water dissociates very slightly into ions in an equilibrium process: 2H2O(l) ⇄ H3O+(aq) + OH−(aq).

Ion-Product Constant for Water (Kw)

𝐾w = [H3O+][OH−] = 1.0 × 10−14 at 25°C. In pure water, [H3O+] = [OH−] = 1.0 × 10−7 𝑀 at 25°C.

Acidic solution

[H3O+] > [OH−]

Neutral solution

[H3O+] = [OH−]

Basic solution

[H3O+] < [OH−]

pH

− log[H3O+]. pH < 7.00 acidic; pH = 7.00 neutral; pH > 7.00 basic.

pOH

− log[OH−]

pKw

pH + pOH = 14.00 at 25°C

Strong Acid

Dissociates completely into ions in water. A dilute solution contains no HA molecules, [HA]eq ≈ 0, and [H3O+] = [A−] ≈ [HA]init.

Types of Strong Acids

HCl, HBr, HI, HNO3, HClO4, and H2SO4. Oxoacids where the number of O atoms exceeds the number of ionizable protons by two or more

Strong Base

Dissociates completely into ions in water. Water-soluble compounds containing O2− or OH− ions, usually those of the most active metals.

Weak Acid

Dissociates slightly to form ions in water, with most HA molecules remaining undissociated.

Le Châtelier’s Principle

When a chemical system at equilibrium is disturbed, it reattains equilibrium by undergoing a net reaction that reduces the effect of the disturbance.

Shift to the Left

Net reaction from product to reactant.

Shift to the Right

Net reaction from reactant to product.

Effect of Concentration Increase

If concentration of component A increases, the system reacts to consume some of it.

Effect of Concentration Decrease

If the concentration of component A decreases, the system reacts to produce some of it.

Inert Gas Addition

Adding an inert gas has no effect on the equilibrium position, as long as the volume does not change.

Change in Volume

Changing the volume of the reaction vessel will cause equilibrium to shift if Δngas ≠ 0.

Effect of Temperature Increase

Adds heat. Results in system shifting in the endothermic (heat-absorbing) direction to consume added heat.

Effect of Temperature Decrease

Removes heat. Results in system shifting in the exothermic (heat-releasing) direction to produce more heat.

The Van't Hoff Equation

ln(𝐾2/𝐾1) = (∆𝐻°𝑅𝑥𝑁 / 𝑅) * (1/𝑇1 − 1/𝑇2)

Catalyst

Speeds up a reaction by lowering its activation energy. Speeds up forward and reverse reactions to the same extent, causing a reaction to reach equilibrium more quickly.

Relationship of thermodynamics and equilibrium dynamics

∆𝐺 = 𝑅𝑇 𝑙𝑛 𝑄 - 𝑅𝑇𝑙𝑛 𝐾

Standard State Conditions

Q = 1; ΔG° = –RT ln K

Reaction Table (ICE Table)

Shows the balanced equation, initial quantities, changes in quantities during the reaction, and equilibrium quantities to determine the equilibrium constant.

Equilibrium Constant (K)

Kc = [Products]/[Reactants]. Ratio of product equilibrium concentrations to reactant equilibrium concentrations

Kc and Kp

Kp = Kc (RT)^Δn
Molarity can be replaced with partial pressures of gases. If the amount of gas does not change, Kp = Kc

If Q < K

The reactants must decrease and the products increase. The reaction proceeds toward the products; reactants → products until equilibrium is reached.

If Q > K

The reactants must increase and the products decrease. The reaction proceeds toward the reactants; reactants ← products until equilibrium is reached.

If Q = K

The system is at equilibrium and no further net change takes place.

Heterogeneous Equilibrium

an equilibrium that involves reactants and products in different phases.

Equilibrium

State where the rates of the forward and reverse reactions are equal.