Acids and Bases

Properties of Acids

• Acids taste sour, but should not be tasted.
• Acids change the color of acid-base indicators.
• They react with active metals, producing hydrogen gas ($H_2$).
• Acids neutralize bases to form salts and water.
• They are also conductive, meaning they can conduct electricity.

Properties of Bases

• Bases taste bitter, but should not be tasted.
• Bases change the color of acid-base indicators.
• They have a slippery feel.
• Bases react with acids to produce salts and water.
• They also conduct electricity.

Arrhenius Definition

• Acids produce hydrogen ions ($H^+$) in aqueous solutions.
• Bases produce hydroxide ions ($OH^−$) in aqueous solutions.

Hydronium Ion ($H_3O^+$)

• Hydrogen ions in water are always solvated
• $H^+ + H_2O \rightleftharpoons H_3O^+$

Brønsted-Lowry Definition

• Acid: a proton ($H^+$) donor
• Base: a proton acceptor

Brønsted-Lowry Acid-Base Reaction Essentials

• Brønsted-Lowry acids require a removable, acidic proton.
• Brønsted-Lowry bases need a pair of nonbonding electrons.

Brønsted-Lowry Acid Example

• Acid donates a proton
• $HF_{(aq)} + H_2O_{(l)} \rightleftharpoons H_3O^+{(aq)} + F^-{(aq)}$
• Here, $HF$ is the acid and $H_2O$ is the base.

Brønsted-Lowry Base Example

• Base accepts a proton
• $NH_{3(aq)} + H_2O_{(l)} \rightleftharpoons NH^+{4(aq)} + OH^-{(aq)}$
• $NH_3$ acts as the base, and $H_2O$ acts as the acid.

Amphoteric Substances Defined

• Substances that can act as either acids or bases are amphoteric.
• Water is the most common example.

Conjugate Acid-Base Pairs Defined

• Conjugate acid is formed when a base accepts a proton.
• Conjugate base is formed when an acid donates a proton.
• $acid \rightleftharpoons conjugate \ base + H^+$
• $base + H^+ \rightleftharpoons conjugate \ acid$

Example of Conjugate Acid-Base Pairs

• In the reaction $HF_{(aq)} + H_2O_{(l)} \rightleftharpoons H_3O^+{(aq)} + F^-{(aq)}$
• $HF$ is the acid and $F^−$ is its conjugate base.
• $H_2O$ is the base and $H_3O^+$ is its conjugate acid.

Acid and Base Strength - Strong Acids

• Strong acids completely dissociate in water, forming extremely weak conjugate bases.
• $HCl_{(aq)} + H_2O_{(l)} \rightarrow H_3O^+{(aq)} + Cl^-{(aq)}$
• $HCl$ is a strong acid, so $Cl^−$ is a weak base.

Acid and Base Strength - Weak Acids

• Weak acids only partially dissociate in water, and their conjugate bases are weak.
• $HF_{(aq)} + H_2O_{(l)} \rightleftharpoons H_3O^+{(aq)} + F^-{(aq)}$
• $HF$ is a weak acid, so $F^−$ is a weak base.

Acid and Base Strength - Strong Bases

• Strong bases dissociate completely in water, their conjugate acids being extremely weak.
• $NaOH_{(s)} \xrightarrow{H_2O} Na^+{(aq)} + OH^-{(aq)}$
• $NaOH$ is a strong base, so $Na^+$ is a weak acid.

Acid and Base Strength - Weak Bases

• Weak bases react partially with water, and their conjugate acids are weak.
• $NH_{3(aq)} + H_2O_{(l)} \rightleftharpoons NH^+{4(aq)} + OH^-{(aq)}$
• $NH_3$ is a weak base, so $NH_4^+$ is a weak acid.

Acid and Base Strength - Equilibrium

• In acid-base reactions, equilibrium favors proton transfer from the stronger acid to the stronger base, forming a weaker acid and base.
• $Acid \ + Base \ \rightleftharpoons Conjugate \ Base \ + Conjugate \ Acid$
• $HF$ is a stronger acid than $H_2O$, thus $F^−$ is a weaker base than $OH^−$.

Autoionization of Water Explained

• Since water is amphoteric, one water molecule can act as an acid and another as a base.
• $H_2O_{(l)} + H_2O_{(l)} \rightleftharpoons H_3O^+{(aq)} + OH^-{(aq)}$

Ion-Product Constant Explained

• The equilibrium expression for water autoionization.
• $Kc = \frac{[H_3O^+][OH^-]}{[H_2O]^2}$
• Can be rearranged to $K_c[H_2O]^2 = [H_3O^+][OH^-]$
• The special equilibrium constant: $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14} \ at \ 25^\circ C$

Defining $pH$

• $pH = -log[H_3O^+]$

$pH$ Values

• Neutral solution: $[H_3O^+] = [OH^-]$, $pH = 7$
• Acidic solution: $[H_3O^+] > [OH^-]$, $pH < 7$
• Basic solution: $[H_3O^+] < [OH^-]$, $pH > 7$

Other "p" Scales

• Can be defined for other equilibrium constants.
• For example, $pKw = -log \ K_w$

Equation for $pK_w$

• $pK_w = pH + pOH = 14.00$

Strong Acids - Memorize These

• Hydrochloric acid ($HCl$)
• Hydrobromic acid ($HBr$)
• Hydroiodic acid ($HI$)
• Sulfuric acid ($H_2SO_4$)
• Nitric acid ($HNO_3$)
• Perchloric acid ($HClO_4$)

Strong Bases

• These are soluble hydroxides.
• Alkali metals (Group 1A) with $OH^−$, such as $NaOH$ and $KOH$.
• Alkaline earth metals (Group 2A) with $OH^−$, such as $Ca(OH)_2$, $Sr(OH)_2$, and $Ba(OH)_2$.

Weak Acids Properties

• Weak acids only partially ionize in solution.
• The acid dissociation constant ($K_a$) is the equilibrium constant for the reaction of a weak acid with water.
• $HA_{(aq)} + H_2O_{(l)} \rightleftharpoons H_3O^+{(aq)} + A^-{(aq)}$

Defining $K_a$

• $K_a = \frac{[H_3O^+][A^-]}{[HA]}$
• The larger the $K_a$, the stronger the acid.

Defining $pK_a$

• $pK_a = -log(K_a)$
• The smaller the $pK_a$, the stronger the acid.

Percent Ionization

• Defined as: $\frac{Concentration \ of \ ionized \ acid}{Original \ concentration \ of \ acid} \times 100$
• For $HF$, percent ionization is $\frac{[H_3O^+]}{[HF]_{initial}} \times 100 = \frac{0.012}{0.20} \times 100 = 6.0 %$

Weak Bases Properties

• Bases react with water to produce hydroxide ions.
• $NH_{3(aq)} + H_2O_{(l)} \rightleftharpoons NH^+{4(aq)} + OH^-{(aq)}$
• The base dissociation constant ($K_b$) is the equilibrium constant for this reaction.

Defining $K_b$

• $K_b = \frac{[NH_4^+][OH^-]}{[NH_3]}$
• The larger the $K_b$, the stronger the base.

Relation Between $K_a$ and $K_b$

• $K_a \times K_b = K_w$
• $pK_a + pK_b = pK_w = 14$

Salts Defined

• Salts are ionic compounds.
• When salts dissolve in water, their ions can act as acids or bases.

Anions Properties

• Anions are conjugate bases of acids.
• Anions of strong acids are not basic.
• Anions of weak acids are weak bases.

Cations Properties

• Cations containing ionizable protons are acidic.
  • Example: $NH_4^+ \rightleftharpoons NH_3 + H^+$
• Metal cations are Lewis acids, attracting water molecules.
• Higher cation charge increases acid strength.

Examples of Salts

• $NaCl$ solution is neutral: $Na^+$ - neutral (cation of a strong base), $Cl^-$ - neutral (anion of a strong acid)
• $NH_4Cl$ solution is acidic: $NH_4^+$ - acidic (cation of a weak base), $Cl^-$ - neutral (anion of a strong acid)
• $NaF$ solution is basic: $Na^+$ - neutral (cation of a strong base), $F^-$ - basic (anion of a weak acid)