Acids and Bases Notes

Acids and Bases

Arrhenius Definition

• Acids produce $H^+$ ions in water.

• Bases produce $OH^-$ ions in water.

• In reality, $H^+$ ions (naked protons) immediately combine with water to form hydronium ions ($H_3O^+$).

  • Not: $HCl \rightarrow H^+ + Cl^-$

  • Actually: $HCl + H<em>2O \rightarrow H</em>3O^+ + Cl^-$

Bronsted-Lowry Definition

• Acids are proton ($H^+$) donors.

• Bases are proton ($H^+$) acceptors.

• Conjugate acid-base pairs differ by one proton ($H^+$).

Conjugate Acid-Base Pairs

• Every acid-base reaction contains two sets of conjugate acid-base pairs.

• A conjugate acid is formed when a proton is transferred to the base.

• A conjugate base is the initial acid minus a proton.

Ionization Equations

• Acids in water:

  • $HNO<em>3 + H</em>2O \rightarrow H<em>3O^+ + NO</em>3^-$

  • $NH<em>4^+ + H</em>2O \rightarrow H<em>3O^+ + NH</em>3$

  • $H<em>2SO</em>4 + H<em>2O \rightarrow H</em>3O^+ + HSO_4^-$

• Bases in water:

  • $F^-(aq) + H_2O(l) \rightarrow HF(aq) + OH^-(aq)$

  • $CH<em>3NH</em>2(aq) + H<em>2O(l) \rightarrow CH</em>3NH_3^+(aq) + OH^-(aq)$

Oxyacids

• Oxyacids have an acidic hydrogen bound directly to oxygen (H-O-X).

• The more electronegative oxygens bound to X, the stronger the acid due to inductive withdrawal of electron density.

  • $H<em>2SO</em>4$ is more acidic than $H<em>2SO</em>3$.

  • $HNO<em>3$ is more acidic than $HNO</em>2$.

Amphoteric Substances

• Amphoteric substances can behave as either an acid or a base.

• Examples: $H<em>2PO</em>4^-$, $HSO_4^-$

Autoionization of Water

• Water can undergo an acid-base reaction with itself due to its amphoteric nature.

• At 25°C, the concentrations of hydrogen and hydroxide ions are equal: $[H^+] = [OH^-] = 1.0 \times 10^{-7} M$.

Strong vs. Weak Acids and Bases

• Strong acids and strong bases completely dissociate into ions.

• Weak acids and weak bases only partially dissociate into ions.

• The reverse reaction is not important for strong acids/bases due to full dissociation.

Concentration vs. Strength

• Concentration is not the same as strength.

• Terms to describe acids/bases: strong, weak, dilute, concentrated.

Acid Strength

• Strong Acid:

  • Equilibrium lies far to the right (e.g., $HNO<em>3 \leftrightarrow H^+ + NO</em>3^-$).

  • Weak H-X bonds, large bond polarity.

  • Yields a weak conjugate base (e.g., $NO_3^−$).

  • Conjugate base has a weak attraction for $H^+$.

• Weak Acid:

  • Equilibrium lies far to the left (e.g., $CH<em>3COOH \leftrightarrow H^+ + CH</em>3COO^-$).

  • Strong H-X bonds, low bond polarity.

  • Yields a relatively strong conjugate base than water (e.g., $CH_3COO^−$).

Acid Dissociation Constant (Ka)

• General form: $HA(aq) \leftrightarrow H^{+1}(aq) + A^{-1}(aq)$

• $K_a = \frac{[H^+][A^-]}{[HA]}$

• For strong acids, $K_a$ is very large.

• For weak acids, $K_a$ is small (often less than one).

Conjugate Strength

• Strong acids have weak conjugate bases and vice versa.

• The reaction proceeds exclusively to the right for strong acids/bases.

Equilibrium Direction

• The stronger acid reacting to form the weaker acid is always favored.

Behavior of Acids in Solution

• Strong acid: Major species are $H_2O$, $H^+$, $A^-$.

• Weak acid: Major species are $H_2O$, $HA$.

Measuring Acid-Base Strength

• Litmus Paper:

  • Red litmus paper turns blue in the presence of a base.

  • Blue litmus paper turns red in the presence of an acid.

• Indicator Paper: Has a color scale to determine a rough pH of the solution.

The pH Scale

• pH is a logarithmic scale based on $[H^+]$.

• $pH = -log[H^+]$ or $pH = -log[H_3O^+]$

• pH of 7 is neutral ($[H^+] = 1.0 \times 10^{-7} M$).

• pH < 7 is acidic ([H^+] > 1.0 \times 10^{-7} M).

• pH > 7 is basic ([H^+] < 1.0 \times 10^{-7} M).

• The pH scale generally ranges from 0-14 but can be negative or greater than 14.

The pOH Scale

• $pOH = -log[OH^-]$

• pOH of 7 is neutral ($[OH^-] = 1.0 \times 10^{-7} M$).

• pOH < 7 is basic ([OH^-] > 1.0 \times 10^{-7} M).

• pOH > 7 is acidic ([OH^-] < 1.0 \times 10^{-7} M).

pH Formulas

• $pH = -log[H^{+1}]$

• $pOH = -log[OH^{-1}]$

• $K_w = [H^{+1}][OH^{-1}] = 1.0 \times 10^{-14}$

• $pH + pOH = 14$

Finding pH and pOH of Strong Acids/Bases

• Strong acids/bases fully dissociate.

• Weak acids/bases require an ICE table.

Undoing pH and pOH

• To find $[H^+]$ or $[OH^-]$ from pH or pOH, take the negative of the inverse log (antilog) of the value.

Weak Acid Equilibria

• $HA(aq) + H<em>2O(l) \rightleftharpoons H</em>3O^+(aq) + A^–(aq)$

• Use the acid dissociation constant ($K_a$) to determine concentrations of $[HA]$, $[H^+]$, and $[A^-]$.

Finding [H3O+] in a Weak Acid Solution

• Use ICE tables to find the equilibrium concentration of $H_3O^+$.

• If it is a strong acid, assume full dissociation; no ICE needed.

Ka from pH

• Use pH to find the $H_3O^+$ concentration at equilibrium, then work backwards through the ICE table to calculate the initial concentration.

Polyprotic Acids

• Polyprotic acids have multiple acidic protons (e.g., $H<em>2SO</em>4$, $H<em>3PO</em>4$, $H<em>2CO</em>3$).

• They ionize in steps, with each successive $H^+$ becoming less acidic.

  • Removing a $H^+$ from an anion is harder than removing it from a neutral compound.

  • The presence of $H^+$ from the first dissociation forces the equilibrium position of the second dissociation to the left (LeChatlier’s principle).

Weak Bases

• Weak base problems are analogous to weak acid problems but produce $OH^-$ instead of $H^+$.

• Many weak bases are Lewis bases that are electron-pair donors through an amine functional group.

Kb

• Base dissociation constant

Neutralization Reactions

• Acid + Base → water + a salt

• Salts: ionic compounds that do not contain $H^+$ ions or $OH^-$ ions.

• Examples:

  • $HCl(aq) + NaOH(aq) \rightarrow H_2O(l) + NaCl(aq)$

  • $H<em>2SO</em>4(aq) + 2LiOH(aq) \rightarrow 2H<em>2O(l) + Li</em>2SO_4(aq)$

Titrations

• In an acid-base titration, an acid or base of known concentration is added to a solution of acid or base of unknown concentration while monitoring the pH with a meter or indicator.

• The point at which equal moles of acid and base have been added is called the equivalence point.

• At the equivalence point, the acid and base are stoichiometrically equivalent.

• Process in which a solution with a known concentration is used to determine an unknown concentration of a second solution.

  • Acid/Base Titration

    • Using an acid or a base of known concentration to determine the unknown concentration of a base or acid.
      An appropriate indicator such as phenolphthalein is used that changes colors at a pH near 7.

Acid/Base Titration Calculations

• $M<em>1V</em>1 = M<em>2V</em>2$

Hydrolysis

• Determining the parent acid and base of a salt and whether it would be acid, basic or neutral.

• Generally:

  • SA + SB = neutral

  • WA + SB = basic

  • SA + WB = acidic

  • WA + WB = varies

Acid-Base Properties of Salts

• Salts that have no effects on pH:

  • The conjugate bases of SA’s and the conjugate acids of SB’s are incredibly weak, so when they dissolve in water they DO NOT ionize water.

  • Examples: KCl, NaCl, NaNO3, KNO3, CsClO4

• Salts that have an effect on pH:

  • The conjugate bases of weak acids and the conjugate acids of weak bases will change the pH of a solution, because they are relatively reactive.

  • Example: NaC2H3O2, KF, NH4NO3, Cs3PO4

• What if the salt is made-up of an acidic cation and a basic anion?

  • Too complicated of for us to get any kind of quantitative value (pH, [H+], [OH-], etc.)

  • We can only get a qualitative answer (acidic or basic) based on comparing the Ka and Kb of the ions.

  • If NH4C3H3O2 is dissolved in water, will the resulting solution be acidic or basic?

    • Ka of HC2H3O2 =1.8x10-5 and Kb of NH3 = 1.8x10-5

  • If NH4CN is dissolved in water, will the resulting solution be acidic or basic?

    • Ka of HCN = 6.2 x 10 -5 and Kb of NH3 =1.8x10 -5