Chemistry 100 Chapter 14 Acids and Bases

Acids and Bases Acids: sour Bases: bitter or salty

Arrhenius definition: Acids and Bases Arrhenius definition: (If H2O is involved.) Acid: produces H3O+ CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq) H3O+ (Hydronium ion): H+(aq) + H2O(l) H3O+(aq) Base: produces OH- H2O NaOH(s) Na+(aq) + OH-(aq) NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

Brønsted and Lowry definition: (If H2O is not involved.) Acids and Bases Brønsted and Lowry definition: (If H2O is not involved.) Acid: donates H+ (proton) a proton donor Base: accepts H+ (proton) a proton acceptor HCl + H2O Cl- + H3O+ acid base Conjugate base Conjugate acid Conjugate acid-base pair Conjugate acid-base pair

Acids and Bases HCl + H2O Cl- + H3O+ Proton (H+) is transferred.

Acids and Bases CH3COOH + NH3 CH3COO- + NH4+ Conjugate Conjugate acid-base pair C6H5OH + H2O C6H5O- + H3O+ acid base Conjugate base Conjugate acid Conjugate acid-base pair Conjugate acid-base pair

Weak acid and base: is partially ionized in aqueous solution. Acids and Bases Weak acid and base: is partially ionized in aqueous solution. produces less H+ and OH- CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq) NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) Strong acid and base: is completely ionized in aqueous solution. produces more H+ and OH- HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq) NaOH(aq) + H2O(l) Na+(aq) + OH-(aq)

Electrolytes + - Electrolyte: conducts an electric current. Na+ Cl- bulb Electrolyte: conducts an electric current. Ionization (Dissociation) NaCl → Na+ + Cl- strong electrolytes: molecules dissociate completely into ions (NaCl). weak electrolytes: molecules dissociate partially into ions (CH3COOH). nonelectrolytes: molecules do not dissociate into ions (DI water).

Acids and Bases Strong acid/base  Strong electrolyte Weak acid/base  Weak electrolyte HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq)

Acids and Bases A strong acid contains a weak conjugate base.

Acids and Bases Monoprotic acids HCl Diprotic acids H2SO4 Triprotic acids H3PO4 Amphiprotic: it can act as either an acid or a base. HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq) NaOH(aq) + H2O(l) Na+(aq) + OH-(aq) base acid

Organic acids: contain carboxyl group (-COOH). Acids and Bases Oxyacids: acidic H is attached to an oxygen atom. H2SO4 H3PO4 HNO3 Organic acids: contain carboxyl group (-COOH). They are usually weak. CH3COOH

Naming binary acids Hydro -ide ion -ic acid Anion : + HF F-: flouride ion Hydroflouric acid HCl Cl-: chloride ion Hydrochloric acid H2S S2-: sulfuride ion Hydrosulfuric acid

Naming ternary acids -ite ion -ous acid Anion: -ate ion -ic acid HNO2 NO2-: Nitrite ion Nitrous acid HNO3 NO3-: Nitrate ion Nitric acid H2CO3 CO32-: carbonate ion carbonic acid H2SO3 SO32-: sulfurite ion sulfurous acid

Ionization constant HA + H2O A- + H3O+ [A-] [H3O+] K = [HA] [H2O] Equilibrium constant K = not for strong acids [HA] [H2O] Ka = K [H2O] = [A-] [H3O+] [HA] Acid ionization constant Ka < 1 - Log Ka = pKa Ka ↑ or pKa ↓ Stronger acid

Ionization of water H2O(l) + H2O(l) ⇌ H3O+ (aq) + OH- (aq) KW = [H3O+] [OH-] = (1×10-7) (1×10-7) [H3O+] [OH-] = 1×10-14 pH + pOH = 14

[H+] and [OH-] [H+] = [OH-] Neutral solution [H+] > [OH-] Acidic solution [H+] < [OH-] Basic solution

pH and pOH pH = - log [H3O+] or -log [H+] pOH = - log [OH-] pH scale: 7 14 Acid Neutral Base [H3O+] ↑ and [OH-] ↓ [H+] = [OH-] [H3O+] ↓ and [OH-] ↑

pH meter and pH indicators

Nature & pH indicators Bigleaf Hydrangea In basic soil (alkaline) In acidic soil

pH of strong acids 0.10 M HCl  pH = ? [H3O+] = [H+] = 0.10 M HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq) 1 mol 1 mol 1 mol 0.10 M 0.10 M 0.10 M [H3O+] = [H+] = 0.10 M pH = -log [H+] pH = -log (0.10) = 1.00

Acid Base Reactions Neutralization: reaction between an acid and a base. Acid + Base Salt + Water KOH(aq) + 2HCl(aq) KCl(aq) + H2O(l) 2NaOH(aq) + H2SO4(aq) Na2SO4(aq) + 2H2O(l) Strong acid reacts with strong base to produce the weaker acid and weaker base. (This is the direction of a reaction)

Titration (Neutralization reaction) B A MB: known VB: known MA: unknown VA: known Equivalence point: Equal amount of acid (H+) and base (OH-) (pH = 7). H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O(l) Acid Base

Practice: Titration (Neutralization reaction) H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O(l) Find the concentration of NaOH solution if 143 mL of this solution is completely neutralized by 0.205 L of 0.150 M H2SO4 solution? 0.150 mol H2SO4 1L H2SO4 solution 1 mol H2SO4 2 mol NaOH × 0.205 L H2SO4 solution × = 0.0615 mol NaOH solution M 0.0615 mol NaOH M = M = = 0.430 M V (L) 0.143 L NaOH

Buffers pH stays constant. Buffer Acid or Base pH stays constant. Buffer A buffer resists changes in pH when limited amounts of acid or base are added.

Buffers Our blood is a buffer solution. pH of blood ≈ 7.4 Acid Acid pH of blood ≈ 7.4 Base Shock Absorber Base

Weak Acid + its Conjugate base (in equilibrium) Buffer Composition Weak Acid + its Conjugate base (in equilibrium) salt of the weak acid CH3COOH + CH3COO-Na+ CH3COOH / CH3COO- Or it can be weak base with it’s conjugate acid.

Buffers pH of blood = between 7.35 and 7.45 Carbonate buffer H2CO3 / HCO3- Phosphate buffer H2PO4- / HPO42- Proteins buffer

How do buffers work? HCO3- + H3O+ → H2CO3 + H2O Carbonate buffer H2CO3 / HCO3- If we eat an acidic food: HCO3- + H3O+ → H2CO3 + H2O H2CO3 + OH- → HCO3- + H2O If we eat a basic food:

Henderson-Hasselbalch equation pH of Buffers HA(aq) A-(aq) + H+(aq) Weak acid Conjugate base [Conjugate Base] pH = pKa + log [Weak Acid] Henderson-Hasselbalch equation [Weak Acid]: concentration of the weak acid [Conjugate Base]: concentration of its conjugate base pKa of the weak acid