Chemistry Reaction Rates and Equilibrium

These flashcards cover key concepts related to reaction rates, equilibrium, and thermodynamics in chemistry.

Endothermic

A process that absorbs heat from its surroundings, often causing a decrease in temperature.

Exothermic

A process that releases heat into its surroundings, often causing an increase in temperature.

Reaction Rate

The speed at which reactants are converted into products in a chemical reaction.

Integrated Rate Law

Mathematical expression that relates the concentration of reactants to time.

Collision Model

Theory that explains how reactions occur through the effective collisions of reactant molecules.

Half-Life

The time required for half of the reactant to be consumed in a reaction.

Zero Order Reaction

A reaction rate that is independent of the concentration of the reactants.

First Order Reaction

A reaction rate that is directly proportional to the concentration of one reactant.

Second Order Reaction

A reaction rate that is proportional to the square of the concentration of one reactant.

Arrhenius Equation

k = Ae^(-Ea/RT), describing how the rate constant (k) depends on temperature and activation energy (Ea).

Rate Determining Step

The slowest step in a reaction mechanism that determines the overall rate of the reaction.

Equilibrium Constant (K)

A value that expresses the ratio of the concentration of products to reactants at equilibrium.

Le Chatelier's Principle

If an external change is applied to a system at equilibrium, the system will adjust to counteract that change. Specific rules include:

1. Change in Concentration: Adding reactants shifts equilibrium to the right (products); adding products shifts to the left (reactants). Removing reactants shifts to the left; removing products shifts to the right.
2. Change in Pressure/Volume: For gaseous reactions, increasing pressure (decreasing volume) shifts equilibrium to the side with fewer moles of gas. Decreasing pressure (increasing volume) shifts to the side with more moles of gas. If moles are equal, pressure change has no effect.
3. Change in Temperature: For an endothermic reaction (\Delta H > 0), increasing temperature shifts equilibrium to the right (products). For an exothermic reaction (\Delta H < 0), increasing temperature shifts equilibrium to the left (reactants). Temperature changes also alter the value of the equilibrium constant (K).

Reaction Quotient (Q)

A measure of the relative amounts of products and reactants present in a reaction at any point in time.

Buffer

A solution that resists changes in pH when small amounts of acid or base are added.

pH Scale

A logarithmic scale used to measure the acidity or basicity of a solution.

Spontaneous Process

A process that occurs without external intervention under a given set of conditions.

Strong Acid

An acid that completely dissociates in water, producing a high concentration of H^+ ions. Common examples include HCl, HBr, HI, HNO3, H2SO4, HClO3, HClO_4.

Weak Acid

An acid that only partially dissociates in water, producing a relatively low concentration of H^+ ions. Most organic acids are weak acids (e.g., CH_3COOH).

Strong Base

A base that completely dissociates in water, producing a high concentration of OH^- ions. Common examples include group 1 hydroxides (LiOH, NaOH, KOH, RbOH, CsOH) and some group 2 hydroxides (Ca(OH)2, Sr(OH)2, Ba(OH)_2).

Solubility Rules

A set of guidelines used to predict whether an ionic compound will dissolve in water. General rules include:

1. Always Soluble: Group 1 metal ions (Li^+, Na^+, K^+, Rb^+, Cs^+), Ammonium (NH4^+), Nitrates (NO3^-), Acetates (CH3COO^-), Chlorates (ClO3^-), Perchlorates (ClO_4^-).
2. Generally Soluble with Exceptions: Halides (Cl^-, Br^-, I^-) are soluble except with Ag^+, Pb^{2+}, Hg2^{2+}. Sulfates (SO4^{2-}) are soluble except with Ca^{2+}, Sr^{2+}, Ba^{2+}, Pb^{2+}, Hg_2^{2+}$.
3. Generally Insoluble with Exceptions: Carbonates (CO3^{2-}), Phosphates (PO4^{3-}), Sulfides (S^{2-}), Oxides (O^{2-}), Hydroxides (OH^-) are insoluble except with Group 1 metals and NH_4^+. Hydroxides are also soluble with $$Ca^{2+}, Sr^{2

Strong Acid

An acid that completely dissociates in water, producing a high concentration of H^+ ions. Common examples include HCl, HBr, HI, HNO3, H2SO4, HClO3, HClO_4.

Weak Acid

An acid that only partially dissociates in water, producing a relatively low concentration of H^+ ions. Most organic acids are weak acids (e.g., CH_3COOH).

Strong Base

A base that completely dissociates in water, producing a high concentration of OH^- ions. Common examples include group 1 hydroxides (LiOH, NaOH, KOH, RbOH, CsOH) and some group 2 hydroxides (Ca(OH)2, Sr(OH)2, Ba(OH)_2).

Solubility Rules

A set of guidelines used to predict whether an ionic compound will dissolve in water. General rules include:

1. Always Soluble: Group 1 metal ions (Li^+, Na^+, K^+, Rb^+, Cs^+), Ammonium (NH4^+), Nitrates (NO3^-), Acetates (CH3COO^-), Chlorates (ClO3^-), Perchlorates (ClO_4^-).
2. Generally Soluble with Exceptions: Halides (Cl^-, Br^-, I^-) are soluble except with Ag^+, Pb^{2+}, Hg2^{2+}. Sulfates (SO4^{2-}) are soluble except with Ca^{2+}, Sr^{2+}, Ba^{2+}, Pb^{2+}, Hg_2^{2+}$.
3. Generally Insoluble with Exceptions: Carbonates (CO3^{2-}), Phosphates (PO4^{3-}), Sulfides (S^{2-}), Oxides (O^{2-}), Hydroxides (OH^-) are insoluble except with Group 1 metals and NH_4^+. Hydroxides are also soluble with $$Ca^{2+}, Sr^{2