Chemistry 102 Ch. 13 Lecture Notes PDF

Preview

Summary

Dr. bethel, Ch 13, an atoms first approach textbook...

Description

Chapter 13: Acids and Bases Section 13.1- The Nature of Acids and Bases -

-

Acids and Bases: A Review - Arrhenius Acid: any substance that when dissolved in water increases concentration of hydrogen ions, H+ - Arrhenius Base: any substance that increases concentration of hydroxide ions, OH- when dissolved in water - Bronsted-Lowry acid: proton (H+) donor - Bronsted-Lowry Base: proton acceptor The Bronsted-Lowry Concept of Acids and Bases - Proton donors may be molecular compounds, cations or anions - HNO3 (aq) + H2O(g) → NO3-(aq) + H3O+ (aq) -

-

NH4+ + H2O → NH3(aq) +H3O+ (aq)

- HCO3-(aq) + H2O(l) → CO3^2-(aq) + H3O+(aq) Proton Acceptors may be molecular compounds, cations or anions - NH3(aq) + H2O(l) → NH4+(aq) + OH-(aq)

- Al(H2O)5(OH)^2+(aq) + H2O(l) → Al(H2O)6^3+(aq) + OH-(aq) Acid In water - HA (aq) + H2O(l) → ← H3O+(aq) + A- (aq) -

Acid + base → ← conjugate acid + conjugate base

-

Loses H+ → conjugate base is everything that remains of the acid molecular after a proton (H+) is lost

-

Gains H+ → conjugate acid is formed when the proton (H+) is transferred to the base

Section 13.2- Acid Strength -

Acid and Base Strength - Strong Acids: dissociate completely in water - Weak Acids: partially dissociate - Strong Bases: dissociate completely in water - Weak Bases: partially dissociate in water - *strong acids and bases go 100% to products, weak acids and bases are mostly still reactants

-

-

-

Strong Acids: - HCl - HBr - HI - HNO3 - HClO4 - H2SO4 Strong Bases: - LiOH - NaOH - KOH - RbOH - Sr(OH)2 - Ba(OH)2 Predicting the Direction of Acid-Base Reactions - According to Bronsted-Lowry theory, all acid-base reactions can be written as equilibria involving the acid and base AND their conjugates - All protons transfer reactions proceed from stronger acid and base to weaker acid and base - Acid + Base → ← Conjugate base of the acid + Conjugate acid of the base When a weak acid is in solution, the products are a strong conjugate acid and base (equilibrium lies to left) When naming acids: - Look at conjugate base → - Ends in -ide: hydro___ic acid - Ends in -ite: ___ous acid -

-

Ex 1. H2PO4+

+

ACID

HCO3 BASE

→←

HPO42CONJ. BASE

+

H2CO3 CONJ. ACID

Ex 2. Write the reaction of perchloric acid with water (identify the stronger acid and weaker acid) HClO4(aq) + H2O(aq) →← ClO4-(aq) + H3O+(aq) ACID STRONGER

BASE STRONGER

CONJ. BASE WEAKER

CONJ. ACID WEAKER

Ex 3. Write the reaction of hydrofluoric acid with water (same thing as above).

HF(aq) + ACID WEAKER

H2O(aq) →← BASE WEAKER

F-(aq) + CONJ. BASE STRONGER

H3O+(aq) CONJ> ACID STRONGER

Ex 4. Write the dissociation expression for the following acids and bases: HCl: HCl(aq) + H2O(l) → ← ClO-(aq) + H3O+(aq) H3PO4 : H3PO4 (aq) + H2O(l) → ← H2PO4-(aq) + H3O+(aq) NH3 : NH3(aq) + H2O(l) → ← NH4+ (aq) + H3O+(aq) - Weak bases tend to gain H+ and have conjugate acid OH-

-

-

Autoionization of Water - The autoionization of water is a reaction by which two water molecules reaction to form H3O+ and OH- K(w) = [H3O+][OH-] = 1.0E-14 @ 25 C - ONLY happens to a very very small extent - Kw is special symbol known as autoionization constant for water Equilibrium Constants for Acids and Bases - The relative strength of an acid or base can also be expressed quantitatively with an equilibrium constant, often called an ionization constant. For the general acid HA, we can write: - HA (aq) + H2O(l) → ← A-(aq) + H3O+(aq) -

-

-

- K(a) = [A-][H3O+(aq)] / [HA] The relative strength of an acid or base can also be expressed quantitatively with an equilibrium constant, often called an ionization constant. For the general base B, we can write: - B(aq) + H2O(l) → ← BH+(aq) + OH-(aq) - K(b) = [BH+][ OH-] / [B] For weak acids and bases - K(a) and K(b) ALWAYS have values that are smaller than one - Acids with a larger K(a) are stronger than ones with a smaller K(a) - Bases with a larger K(b) are stronger than ones with smaller K(b) The weaker acid, the stronger its conjugate base: the smaller the value of K(a), the

-

larger the value of K(b) Aqueous acids that are stronger than H3O+ are completely ionized - Their conjugate based (such as NO3-) do not produce meaningful concentrations of OH- ions, their K(b) values are NEGLIGIBLE - Same is true for strong bases and their conjugate acids

Ex 1. How does the ionization constant for H2S (an acid) relate to its conjugate base? K(a) H2S(aq) + H2O(l) → ← HS-(aq) + H3O+(aq) HS-(aq) + H2O(l) → ← H2S(aq) + H3O+(aq)

K(b)

ADD TOGETHER → 2H2O (l) → ← H3O+(aq) + OH-(aq) Ex 2. THe K(a) of benzoic acid (C6H5COOH) is 6.3E-5. Write the reaction of benzoate (its conjugate base) with water. What is the K(b) of benzoate? - K(a) * K(b) = K(w) C6H5COO- (aq) + H2O(l) → ← C6H5COOH(aq) +OH-(aq) Kb = Kw/Ka = (1.0E-14) / (6.3E-5) = 1.6E-10

Section 13.3- The pH Scale -

-

pH scale basics - The pH of a solution is defined as the negative of the base (10) logarithm (log) of the hydronium ion concentration - pH = -log[H+] - In a similar way, we can define the pOH of a solution as the negative of the base (10) logarithm of the hydroxide ion concentration - pOH = -log[OH-] - The concentration of acid, [H3O+] is found by taking the antilog of the solutions pH - In a similar way [OH-] can be found from: - Once [H3O+] is known, [OH-] can be found from: - VICE VERSA Relative Acid Strength - We know that a weak acid with a larger K(a) value is stronger than a week acid with a smaller K(a)

Ex 1. What is the pH of a 4.2E-3 M HCl solution?

HCl(aq) + H2O(l) → ← H3O+(aq) + Cl-(aq) pH = -log(4.2E-3 M) = 2.38 Ex 2. The pH of bleach is 12.32 at 25 C. What is the hydroxide ion concentration? What is the hydrogen ion concentration. - pH = -log (h3o+) - H3o+ = 10-pH 10-12.32 = 4.8E-13 14.00-12.32 = pOH= 1.68 [OH] = 10-1.68 = 2.1E-2 M

Section 13.4- Calculating the pH of Strong Acid Solutions Ex 1. If the pOH of Ba(OH)2 , solution is 1.54, what is the concentration of Ba(OH)2? - pOH = 1.54 Ex. 2. What is the pH of dilute muriatic acid (a solution that is 5.0% HCl)? You should use 1.00g h2o = 1.00 mL - 5.00 % HCl 5.00 g HCl/ (5.00 g HCl + 95.00 g h2O) 0.137 mol HCl / 0.09500 K soln) = 1.44 M pH = -log(1.44) = -0.159

Section 13.5- Calculating the pH of Weak Acid Solutions -

-

Solving Weak Acid Equilibrium Problems 1. Write out the balanced chemical equations for the reactions producing H+ 2. Using the values of the equilibrium constants for the reactions you have written, decide which equilibrium will dominate in producing H+ 3. Write the equilibrium expression for the dominant equilibrium 4. List the initial concentrations of the species participating in the dominant equilibrium 5. Define the change needed to achieve equilibrium: define x 6. Write the equilibrium concentrations in terms of x 7. Substitue the equilibrium concentrations into the equilibrium expression 8. CHeck to see if you can take a shortcut (K is usually very small) 9. Calculate [H+] and pH 5% Rule - x / [acid or base] * 100%...