Strong Acids and Bases

Identifying Acids and Bases

• Acids: Typically have hydrogen in front of their chemical formula.

  • Examples: HCl (hydrochloric acid), HF (hydrofluoric acid), HC2H3O2 (acetic acid).

• Bases: Typically have a hydroxide ion (OH^-).

  • Examples: NaOH (sodium hydroxide), KOH (potassium hydroxide).

• Exceptions:

  • Hydrogen next to a metal (e.g., NaH - sodium hydride) is a base.

  • Hydrogen attached to a nonmetal is typically an acid.

• Charge:

  • Positive hydrogen indicates an acid.

  • Negative hydrogen indicates a base.

  • Acids tend to be positively charged, bases negatively charged.

Arrhenius Definition

• Acids: Substances that release hydrogen ions (H^+) into solution.

  • Hydrogen ions in water are equivalent to hydronium ions (H_3O^ +) because H^+ doesn't exist by itself in water; it bonds to water.

• Bases: Substances that release hydroxide ions (OH^-) into solution.

Bronsted-Lowry Definition

• Acids: Proton donors.

• Bases: Proton acceptors.

• Example 1: Hydrochloric Acid in Water

  • Reaction: HCl + H2O \rightarrow Cl^- + H3O^ +

  • HCl is the Bronsted-Lowry acid (proton donor).

  • H_2O is the Bronsted-Lowry base (proton acceptor).

  • Cl^- is the conjugate base (formed after HCl donates a proton).

  • H3O^ + is the conjugate acid (formed after H2O accepts a proton).

• Example 2: Ammonia in Water

  • Reaction: NH3 + H2O \rightarrow NH_4^ + + OH^-

  • NH3 is the base (proton acceptor, turns into NH4^ +).

  • H_2O is the acid (proton donor, turns into OH^-).

  • NH_4^ + is the conjugate acid.

  • OH^- is the conjugate base.

Writing Conjugate Acids and Bases

• Conjugate Acid: Add H^+ and increase the charge by 1.

  • Example: Water (H2O) becomes hydronium (H3O^ +).

• Conjugate Base: Remove H^+ and decrease the charge by 1.

  • Example: Water (H_2O) becomes hydroxide (OH^-).

• Examples:

  • Ammonia (NH_3):

    • Conjugate acid: NH_4^ +

    • Conjugate base: NH_2^-

  • Dihydrogen Phosphate (H2PO4^-):

    • Conjugate acid: H3PO4

    • Conjugate base: HPO_4^{2-}

pH Scale

• Typically ranges from 0 to 14, but can go beyond these numbers.

• pH = 7: Neutral solution.

• pH < 7: Acidic solution (e.g., pH = -2 is very acidic).

• pH > 7: Basic solution.

pH Calculations

• pH = -log[H_3O^+] (negative log of hydronium ion concentration).

• pOH = -log[OH^-] (negative log of hydroxide ion concentration).

• pH + pOH = 14 (at 25 degrees Celsius).

• [H_3O^+] = 10^{-pH} (hydronium ion concentration).

• [OH^-] = 10^{-pOH} (hydroxide ion concentration).

Strong vs Weak Acids

• Strong Acids: Ionize completely in solution.

  • Form strong electrolytes (conduct electricity well).

  • Ionize almost 100%.

• Weak Acids: Partially ionize in solution (less than 5%).

  • Form weak electrolytes.

Common Strong Acids

• HCl (hydrochloric acid)

• HBr (hydrobromic acid)

• HI (hydroiodic acid)

• HNO_3 (nitric acid)

• H2SO4 (sulfuric acid)

• HClO_4 (perchloric acid)

• Note: HF (hydrofluoric acid) is a weak acid, not a strong acid.

Weak Acids

• NH_4^ + (ammonium ion)

• HC2H3O_2 (acetic acid)

• Cyanic acid

• Nitrous acid

• Sulfurous acid

Oxyacids

• For oxyacids, the acid with more oxygen atoms is more acidic.

  • Sulfuric acid (H2SO4) is stronger than sulfurous acid (H2SO3).

  • Nitric acid (HNO3) is stronger than nitrous acid (HNO2).

  • Perchloric acid (HClO4) > chloric acid (HClO3) > chlorous acid (HClO_2) > hypochlorous acid (HClO).

• This trend does not apply to acids without oxygen (e.g., HCl).

Chemical Reactions with Strong and Weak Acids

• Strong Acids: Use a single arrow to show complete ionization.

  • Example: HCl + H2O \rightarrow Cl^- + H3O^ +

• Weak Acids: Use a double arrow to show partial ionization and equilibrium.

  • Example: HF + H2O \rightleftharpoons F^- + H3O^ +

Strong vs Weak Bases

• Strong Bases: Soluble ionic compounds that ionize completely.

  • Examples: KOH (potassium hydroxide), NaOH (sodium hydroxide), Ba(OH)_2 (barium hydroxide).

  • Form strong electrolytes.

• Weak Bases: Insoluble compounds that ionize partially (less than 1%).

  • Example: Al(OH)_3 (aluminum hydroxide).

  • Ammonia (NH3) and conjugate bases of weak acids (e.g., fluoride, nitrite, acetate, HSO3^-).

Other Strong Bases

• Oxide (O^{2-}): A stronger base than hydroxide.

  • O^{2-} + H_2O \rightarrow 2OH^-

• Hydride (H^-): A strong base that produces hydrogen gas and hydroxide ions when reacting with water.

  • H^- + H2O \rightarrow H2 + OH^-

Base Strength Comparison

• O^{2-} > OH^- > H2O > H3O^+

• Oxide is the most basic, hydronium is the most acidic.

• Hydride is a stronger base than hydroxide.

Reaction Mechanisms

• Oxide and Water: Oxide uses a lone pair to take a hydrogen from water, forming two hydroxide ions.

• Hydride and Water: Hydride (with a negative charge) is attracted to the partially positive hydrogen in water, forming hydrogen gas and hydroxide ions.

Properties of Acids and Bases

• Taste:

  • Acids taste sour (e.g., lemons).

  • Bases taste bitter.

• Feel:

  • Bases feel slippery.

• Indicators:

  • Acids turn blue litmus paper red.

  • Bases turn red litmus paper blue.

    

pH of Solutions

• Acidic: pH < 7

• Neutral: pH = 7

• Basic: pH > 7

Electrical Conductivity

• Strong acids and strong bases are strong electrolytes; they conduct electricity very well due to complete ionization.

• Weak acids and weak bases are weak electrolytes; they conduct a small amount of electricity due to partial ionization.

• Electrical conductivity of a strong acid solution > electrical conductivity of a weak acid solution.

Reactions with Active Metals

• Acids react with active metals to produce hydrogen gas.

  • Example: Zn + 2HCl \rightarrow ZnCl2 + H2

• Active metals: Zinc, aluminum, iron, nickel, sodium (but sodium is too reactive with water).

• Inactive metals (copper, silver, gold) do not react with acids to produce hydrogen gas.

Definitions of Acids and Bases (Recap)

• Arrhenius:

  • Acids release H^+ ions in solution.

  • Bases release OH^- ions in solution.

• Bronsted-Lowry:

  • Acids are proton donors.

  • Bases are proton acceptors.

• Lewis:

  • Acids are electron pair acceptors.

  • Bases are electron pair donors.

Acid Dissociation Constant (Ka)

• Example: Hydrofluoric Acid in Water

  • Reaction: HF(aq) + H2O(l) \rightleftharpoons H3O^+(aq) + F^-(aq)

  • Ka is the acid dissociation constant.

  • Ka = \frac{[H_3O^+][F^-]}{[HF]} (products over reactants).

  • Liquids and solids are not included in the equilibrium expression.

• As Ka increases, the strength of the acid increases.

• As Ka increases, the pKa value decreases.

• Strong acids have large Ka values, small pKa values.

• pKa = -log(Ka)

Base Dissociation Constant (Kb)

• Example: Ammonia in Water

  • Reaction: NH3(aq) + H2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)

• Kb is the base dissociation constant.

• Kb = \frac{[NH4^+][OH^-]}{[NH3]}

• pKb = -log(Kb)

Amphoteric Substances

• Substances that can act as both an acid and a base.

• Example: Water (H_2O)

• Example: Dihydrogen Phosphate (H2PO4^-)

  • As a base: H2PO4^- + HF \rightarrow H3PO4 + F^-

  • As an acid: H2PO4^- + NH3 \rightarrow NH4^+ + HPO_4^{2-}

• When an acid loses a hydrogen, it forms the conjugate base.

• When a base gains a hydrogen, it forms the conjugate acid.

Autoionization of Water

• Water reacts with itself.

• Reaction: 2H2O(l) \rightleftharpoons H3O^+(aq) + OH^-(aq)

• Kw = [H_3O^+][OH^-] (autoionization constant for water).

• Kw is temperature-dependent (increases with temperature).

• At 25 degrees Celsius, Kw = 1 \times 10^{-14}.

• [H_3O^+][OH^-] = 1 \times 10^{-14}

• pH + pOH = 14

• pKa + pKb = 14

• Ka * Kb = Kw = 1 \times 10^{-14}

Practice Problems

Problem 1

The H_3O^+ concentration in a solution is 4 \times 10^{-3}. What is the pH of the solution?

pH = -log[H_3O^+]

pH = -log(4 \times 10^{-3}) = 2.3979 \approx 2.4

Problem 2

The OH^- concentration in a solution is 5.3 \times 10^{-4}. What is the pOH of the solution?
pOH = -log[OH^-]
pOH = -log(5.3 \times 10^{-4}) = 3.2757 \approx 3.28

Problem 3

We're given the hydronium ion concentration, and we want to calculate the hydroxide ion concentration.

[H_3O^+][OH^-] = Kw

[OH^-] = Kw/[H_3O^+]

[OH^-] = (1 \times 10^{-14}) / (2.5 \times 10^{-5}) = 4 \times 10^{-10}

Problem 4

If the Ka of acetic acid (HC2H3O_2) is 1.8 \times 10^{-5}, what is the pKa of the acid?

pKa = -log(Ka)

pKa = -log(1.8 \times 10^{-5}) = 4.745

Problem 5

Which of the following statements is false?

• Bases taste bitter and feel slippery. (True)

• Acids taste sour and acids react with active metals to produce hydrogen gas. (True)

• HCl is a strong electrolyte. (True)

• Acids turn red litmus paper blue. (False)

• Sodium hydroxide is a strong base, etc. (True)

Problem 6

Which of the following solutions will have the highest pH? (Assuming the same concentration for each).

• HBr (strong acid)

• HF (weak acid)

• NaCl (neutral salt, pH = 7)

• NH_3 (weak base)

• KOH (strong base)

The higher the pH value, the stronger the base.

Problem 7

The Ka value for HF and acetic acid are 7.2 \times 10^{-4} and 1.8 \times 10^{-5}, respectively. Which acid is stronger: hydrofluoric acid or acetic acid?
As the Ka value increases, the strength of the acid increases. Because 7.2 \times 10^{-4} is larger than 1.8 \times 10^{-5}, hydrofluoric acid is the stronger acid.

Problem 8

Match each term with the correct letter.

1. Arrhenius acid (releases H+ ions in solution)

2. Arrhenius base (releases hydroxide ions in solution)

3. Bronsted-Lowry acid (Proton donor)

4. Bronsted-Lowry base (Proton acceptor)

5. Lewis acid (electron pair acceptor)

6. Lewis base (electron pair donor)