Equilibrium & Acids and Bases

Equilibrium & Acids and Bases

Unit Overview

Unit 11 covers equilibrium, acids, and bases, including chapters 14, 15, and 18.

Reversible Reactions

Forward reaction: Reactants → Products
Reverse reaction: Reactants ← Products
Reversible Reaction: Reactants ⇋ Products

Chemical Equilibrium

Chemical equilibrium is the state where the concentration of reactants and products remains constant.
The rate at which reactants are formed equals the rate at which they are consumed.

Equilibrium State

Equilibrium is reached when the concentrations of reactants and products remain constant over time.
The ratio of reactants to products will be the same no matter the starting point. For example:
$H<em>2 (g) + I</em>2 (g) \rightarrow 2 HI (g)$

Le Chatelier's Principle

If a system at equilibrium experiences a change in conditions, the equilibrium position will shift to counteract that change.

Factors Affecting Equilibrium

Changes in concentration
Changes in temperature
Changes in pressure

Changes in Concentration

If the concentration of a substance in a reversible reaction is increased, the equilibrium will shift away from that substance.
Example: $2 NO<em>2 (g) \leftrightarrow N</em>2O_4 (g)$
- Adding more $NO_2$ shifts the reaction to the right, producing more products.
- Adding more $N<em>2O</em>4$ shifts the reaction to the left, producing more reactants.

Changes in Pressure

If pressure is increased, the reaction will shift towards the side with fewer moles of gas.
If pressure is decreased, the reaction will shift towards the side with more moles of gas.
Example: $2 NO<em>2 (g) \leftrightarrow N</em>2O_4 (g)$
- Increasing pressure shifts the reaction to the right, making more products.
- Decreasing pressure shifts the reaction to the left, making more reactants.

Changes in Temperature

Exothermic Reactions: Consider heat as a product.
Endothermic Reactions: Consider heat as a reactant.

Exothermic Reactions

Example: $H<em>2 (g) + I</em>2 (g) \leftrightarrow 2 HI (g) + heat$
- Increasing temperature shifts the reaction to the left, producing more reactants.
- Decreasing temperature shifts the reaction to the right, producing more products.

Endothermic Reactions

Example: $Heat + NH<em>4Cl (s) \leftrightarrow NH</em>3 (g) + HCl (g)$
- Increasing temperature shifts the reaction to the right, producing more products.
- Decreasing temperature shifts the reaction to the left, producing more reactants.

Application: Producing Ammonia

Producing Ammonia ( $NH_3$ ) for explosives and fertilizer.
Reaction:
$N<em>2 (g) + 3 H</em>2 (g) \leftrightarrow 2 NH_3 (g) + heat$
Conditions that increase the amount of ammonia produced: High pressure and low temperature.

Reaction Rates

Reaction rates measure how much of the reactants are used up or products are produced per unit of time.

Requirements for a Reaction

Molecules need to collide with one another.
Molecules need to collide with enough energy to break current bonds and form new bonds.
Reactions occur faster when particles collide more frequently and with more force.

Activation Energy

Activation energy is the minimum energy colliding particles must have in order to react.
It represents the energy difference between reactants and the transition state.

Gibbs Free Energy

A diagram illustrating Gibbs Free Energy shows the energy required for reactants to reach the transition state and form products.

Ways to Change Reaction Rates

Temperature
Concentration
Particle Size
Catalyst

Temperature

Higher temperature -> Particles move faster -> Harder collisions -> Increased number of collisions -> Increased reaction rates

Concentration

Increased concentration increases the frequency of collisions, leading to a greater rate of reaction.

Particle Size

Smaller particle size increases the surface area for the reaction.
More surface area allows for more frequent and effective collisions.

Catalyst

A catalyst is a substance that increases the rate of reaction.
It is not consumed in the reaction.
It lowers the activation energy, making it easier for the reaction to occur.

Self-Ionization of Water

Water reacts with itself in a reversible reaction to produce hydronium ( $H_3O^+$ ) and hydroxide ( $OH^-$ ) ions.
Reaction: $H<em>2O + H</em>2O \rightleftharpoons H_3O^+ + OH^-$
Occurs at very low levels.
Pure water at 25°C is a neutral solution.
$[H^+] = [OH^-] = 1.0 \times 10^{-7} M$
Ion product constant ( $K<em>w$ ): $K</em>w = [H^+] \times [OH^-] = 1.0 \times 10^{-14} M$

Acidic Solutions

[H^+] > [OH^-]
[H^+] > 1.0 \times 10^{-7} M
[OH^-] < 1.0 \times 10^{-7} M

Properties of Acids

Taste sour.
Sting cuts.
Produce hydrogen gas when reacting with metal.
Have a pH < 7.

Basic Solutions

[OH^-] > [H^+]
[H^+] < 1.0 \times 10^{-7} M
[OH^-] > 1.0 \times 10^{-7} M

Properties of Bases

Taste bitter.
Feel smooth/slippery.
Do not react with metals.
Have a pH > 7.

Acid-Base Definitions

Arrhenius Acid

Hydrogen-containing molecules that produce $H^+$ when dissolved in water.

Arrhenius Base

Compounds that ionize to produce hydroxide ( $OH^-$ ) in water.

Bronsted-Lowry Acid

Molecules that donate a hydrogen ion ( $H^+$

Bronsted-Lowry Base

Molecules that accept a hydrogen ion.
Example: $NH<em>3 + H</em>2O \rightleftharpoons NH_4^+ + OH^-$

Conjugate Acids/Bases

Conjugate acid = ion or molecule formed when a base gains a hydrogen ion
Conjugate base = ion or molecule that remains after an acid loses a hydrogen ion
Example:
$NH<em>3(aq) + H</em>2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)$

Examples of Identifying Acids, Bases, Conjugate Acids, and Conjugate Bases

$HClO<em>4(aq) + H</em>2O(l) \rightleftharpoons H<em>3O^+(aq) + ClO</em>4^–(aq)$
$H<em>2SO</em>3(aq) + H<em>2O(l) \rightleftharpoons H</em>3O^+(aq) + HSO_3^–(aq)$
$HC<em>2H</em>3O<em>2(aq) + H</em>2O(l) \rightleftharpoons H<em>3O^+(aq) + C</em>2H<em>3O</em>2^–(aq)$
$H<em>2S(g) + H</em>2O(l) \rightleftharpoons H_3O^+(aq) + HS^–(aq)$
$HSO<em>3^–(aq) + H</em>2O(l) \rightleftharpoons H<em>3O^+(aq) + SO</em>3^{2–}(aq)$

Amphoteric Substances

Substances that can act as either an acid or a base.
Example: $H<em>2CO</em>3 \rightleftharpoons HCO<em>3^- \rightleftharpoons CO</em>3^{2-}$

Strong vs. Concentrated

Strong: Dissociates completely
Weak: Does not completely dissociate
Concentrated: Lots of compound in solution
Dilute: Not a lot of the compound in solution
Language matters when describing acids and bases.

The pH Scale

Measure of the $[H_3O^+]$ in the solution, or how acidic the solution is
Ranges from 0 to 14, with 7 being neutral.
pH < 7 is acidic
pH > 7 is basic (alkaline)
Each step of the pH scale represents a change in concentration by a factor of 10.

pH Scale and Ion Concentrations

In a neutral solution: $[H_3O^+] = [OH^-]$
In an acidic solution: [H_3O^+] > [OH^-]
In a basic solution: [H_3O^+] < [OH^-]

pH Calculations

$pH = -log[H^+]$
$[H^+] = 10^{-pH}$
$pOH = -log[OH^-]$
$[OH^-] = 10^{-pOH}$
$pH + pOH = 14$
$[OH^-][H^+] = 1 \times 10^{-14}$

pH Calculation Formulas

Calculating pH from $[H^+]$ and $[OH^-]$

Examples

What is the pH of a solution with $OH^- = 2.5 \times 10^{-6}$ ?
The pH of blood is 7.4. What is the hydronium concentration in blood?

Neutralization Reaction

Acids and bases react together and neutralize each other.
The result is a salt (ionic compound) and water.
Acid + Base → Salt + Water

Examples

Predict the products of these neutralization reactions
- $HCl (aq) + NaOH (aq) \rightarrow$
- $HF + Mg(OH)_2\rightarrow$
- $H<em>2SO</em>4 + LiOH \rightarrow$

Acid/Base Titrations

A carefully controlled neutralization reaction used to measure the concentration of an acid or a base.
A solution of known concentration (standard) is added to a solution with unknown concentration until the equivalence point is reached.

Titration Curves

Graphs that show the pH as base is added to the solution.
At the equivalence point, the moles of $H^+$ from the acid is equal to the moles of $OH^-$ from the base.

Solving Titration Calculation Problems

Figure out how many moles of the standard were used.
$Molarity = \frac{mol}{L}$
At the equivalence point, moles of $H^+$ = moles of $OH^-$ .
Calculate the molarity of the unknown.
$Molarity = \frac{mol}{L}$

Examples

In a titration, 25 ml of 0.16 M NaOH was required to neutralize a 50 ml sample of HCl. What was the molarity of the HCl.
In a titration, 27 ml of 0.45 M $Sr(OH)<em>2$ is required to neutralize 18 ml of $HNO</em>3$ . Write a balanced chemical equation for the neutralization reaction and determine the molarity of the nitric acid.