Presentation on theme: "Acid-Base Equilibria BLB 12 th Chapter 16. Expectations Distinguish between acids and bases Definitions & properties Know common strong and weak examples."— Presentation transcript:
Acid-Base Equilibria BLB 12 th Chapter 16
Expectations Distinguish between acids and bases Definitions & properties Know common strong and weak examples Calculate pH for strong and weak systems Write chemical reactions of acids and bases Predict relative acid-base strength
Examples of acids & bases
AcidsBases Sour (like vinegar)Bitter and slippery (like soap) React with bases to neutralize them and form salts React with acids to neutralize them and form salts Change indicator colors in opposite direction from base (e.g. litmus blue to red) Change indicator colors in opposite direction from acid (e.g. litmus red to blue) Aqueous solutions conduct electricity Liberate hydrogen in reactions with active metals React in aqueous solution with salts of heavy metals to form insoluble hydroxides or oxides
16.1 Acids & Bases: A Brief Review Arrhenius Definitions Acid – a substance that produces hydrogen ions (H + ) in water HA → H + + A - Base – a substance that produces hydroxide ions (OH - ) in water BOH → B + + OH -
16.2 Brønsted-Lowry Acids & Bases H + (proton) in water: H + + H 2 O → H 3 O + hydronium ion Hydronium ion can hydrogen bond with more water molecules to form large clusters of hydrated hydronium ions. H + and H 3 O + are used interchangeably.
16.2 Brønsted-Lowry Acids & Bases Brønsted-Lowry definitions acid – hydrogen ion (or proton) donor Neutral (HNO 3 ), anionic (HCO 3 - ), cationic (NH 4 + ) Must have a removable (acidic) proton base – hydrogen ion (or proton) acceptor Neutral (NH 3 ), anionic (HCO 3 -, CO 3 2- ) Must have a lone pair of electrons
amphiprotic – capable of behaving as a Brønsted acid and Brønsted base amphoteric – capable of behaving as a Lewis acid and Brønsted base (17.5) Neutralization reaction in which mol acid = mol base acid(aq) + base(aq) → salt(aq) + water(l) HCl(aq) + NaOH(aq) → NaCl(aq) + H 2 O(l)
Conjugate acid/base pairs – reactant and product that differ by a single proton HA(aq) + H 2 O(l) → H 3 O + (aq) + A - (aq) acid + base conj. acid + conj. base HCl(aq) + NH 3 (aq) → NH 4 + (aq) + OH - (aq)
Relative Strengths of Acids and Bases Strength is a measure of the ability of an acid (or base) to donate (or accepts) a H +. Stronger acids donate H + more readily. Completely dissociate in water Conjugate bases have negligible tendency to accept protons; neutral. Weaker acids donate H + less readily. Partially dissociate and establish equilibrium Conjugate bases have some tendency to accept protons. The stronger an acid, the weaker its conjugate base and vice versa.
HA(aq) + H 2 O(l) ⇌ H 3 O + (aq) + A - (aq) HA(aq) + H 2 O(l) → H 3 O + (aq) + A - (aq)
Acid/base reactions proceed from the stronger acid-base pair to the weaker acid- base pair. Common strong acids (p. 664): HClO 4, HClO 3, H 2 SO 4, HI, HBr, HCl, HNO 3 Monoprotic acid – capable of donating only one H + Polyprotic acid – capable of donating more than one H + Common strong bases (p. 665): M(OH) n, where M = Group I (n=1) & heavier Group II (n=2) metals
16.3 The Autoionization of Water H 2 O(l) + H 2 O(l) ⇌ H 3 O + (aq) + OH - (aq) H 2 O(l) ⇌ H + (aq) + OH - (aq) K w = [H 3 O + ][OH - ] = [H + ][OH - ] = 1.0 x °C) K w – ion-product constant (or dissociation constant) Pure water is neutral. Thus, [H 3 O + ] = [OH - ] = 1.0 x °C
For an aqueous solution: 16.3 The Autoionization of Water
Working with K w
16.4 The pH Scale pH represents a solution’s acidity 25°C). 0 ← 7 → 14 acidic neutral basic See Table 16.1, p. 661 for summary. See Figure 16.5, p. 663 for examples. pH = −log[H 3 O + ] = −log[H + ] [H 3 O + ] = 10 -pH pOH = −log[OH - ] pH + pOH = 14 [OH - ] = 10 -pOH
More common chemicals ChemicalpH Basic Windex10.57 Bleach9.58 Neutral Tap water*7.46 Acidic Alka Seltzer (in tap water)6.43 Distilled water**6.37 Flat Coke2.60 Toilet bowl cleaner M HCl−0.29 *CaCO 3 CO H 2 O ⇌ HCO OH - **CO 2 + H 2 O → H 2 CO 3
More about pH pH does not necessarily indicate strength. Measuring pH pH meters – measures exact pH based on electrochemistry Acid-base indicators – estimates pH based on the appearance of color
16.5 Strong Acids and Bases Strong acids & bases completely dissociate. [HA] 0 = [H 3 O + ] → pH [MOH] 0 = [OH - ] → pOH → pH 2[M(OH) 2 ] 0 = [OH - ] → pOH → pH H 3 O + is the strongest acid that can exist in water. (produced by all acids in water) OH - is the strongest base that can exist in water. (produced by all bases in water)
pH problems End Test #1 material
16.6 Weak Acids & 16.7 Weak Bases Weak acids & bases do not completely dissociate. Weak acids establish an equilibrium in aqueous solution. HA(aq) + H 2 O(l) ⇌ H 3 O + (aq) + A - (aq) HA(aq) ⇌ H + (aq) + A - (aq) They do not readily donate or accept H + ’s. [HA] 0 ≠ [H 3 O + ] [MOH] 0 ≠ [OH - ]
HA(aq) + H 2 O(l) ⇌ H 3 O + (aq) + A - (aq) HA(aq) ⇌ H + (aq) + A - (aq) K a ↑ acid strength ↑ For polyprotic acids: K a1 >> K a2 >> K a3 pK a = −log[K a ] pK a ↑ acid strength↓ Weak Acids & Acid-dissociation Constant
From p more in Appendix D, p. 1062
Weak Bases & Base-dissociation Constant Weak bases establish an equilibrium in aqueous solution. B(aq) + H 2 O(l) ⇌ BH + (aq) + OH - (aq) K b ↑ base strength ↑ pK b = −log[K b ] pK b ↑ base strength↓
Weak acid/base Problems 1) K a (or K b ) from equilibrium pH 2) pH from K a (or K b ) 1. Identify as weak acid or base. 2. Write the chemical equilibrium. 3. Write the equilibrium constant expression. 4. Set up concentration table. (Ch. 15.5) 5. Solve for x. 6. Check with 5% rule. If greater than 5%, use quadratic equation. (type 2 only) 7. Complete problem.
The pH of a 0.10 M solution of propionic acid (CH 3 CH 2 CO 2 H) is Calculate the K a for propionic acid.
Calculate the pH of a 1.0 M HF solution.
Calculate the pH of a M HF solution.
Calculate the pH of a 0.20 M solution of triethylamine N(CH 2 CH 3 ) 3.
16.8 Relationship between K a and K b For a conjugate acid/base pair: K a x K b = K w (derivation p. 679) Thus, at 25°C, K a x K b = 1.0 × And, pK a + pK b = pK w = 14.00
16.9 Acid-Base Properties of Salt Solutions Salt – ionic compound Salts dissolve in water to produce ions. Ions can also affect the pH. Hydrolysis – reaction between an ion and water to produce H 3 O + or OH - F - (aq) + H 2 O(l) ⇌ HF(aq) + OH - (aq) NH 4 + (aq) + H 2 O(l) ⇌ H 3 O + (aq) + NH 3 (aq)
Which ions will undergo hydrolysis, i.e. react with water and affect the pH of the solution? Anion: Conjugate base of a weak acid ► basic Conjugate base of a monoprotic strong acid ► neutral Cation: Conjugate acid of a weak base ► acidic Group I & II metal ions ► neutral (exceptions Be 2+ and Mg 2+ ► acidic) Other metal ions ► acidic See p. 683 for summary of combined effect.
Effect on cations on solution pH
Cation + Anion ►Acidic, basic, or neutral?
Calculate the pH of a 0.15 M NaC 2 H 3 O 2, sodium acetate, solution.
16.10 Acid-Base Behavior and Chemical Structure Binary Acids (HX) As bond H−X bond strength increases, acid strength decreases. The greater the stability of the conjugate base, X -, the stronger the acid. Group: size of X ↑, bond strength ↓, acid strength ↑ Period: electronegativity of X ↑, acid strength↑
Oxyacids – acidic H attached to an oxygen atom Same # of OH groups and O atoms: central atom electronegativity ↑, acid strength ↑ HClO > HBrO > HIO Same central atom, Y: # O atoms ↑, acid strength ↑ HClO 4 > HClO 3 > HClO 2 > HClO Carboxylic acids – contain −COOH or CO 2 H # electronegative atoms ↑, acid strength ↑
16.11 Lewis Acids and Bases Lewis acid – electron-pair acceptor e - -poor compounds Metal ions Lewis base – electron-pair donor Amines, NR 3 Ligands (see chapter 23.2) Every Brønsted acid/base is a Lewis acid/base, but not vice versa.
16.11 Lewis Acids and Bases
Lewis acid & base examples
Amphoterism – capable of acting as a Bronsted base and a Lewis acid (See p. 733 for more information.)