Acids and Bases Chapter 13

Some Properties of Acids Produce H+ (as H3O+) ions in water (the hydronium ion is a hydrogen ion attached to a water molecule) Taste sour Corrode metals Electrolytes React with bases to form a salt and water pH is less than 7 Turns blue litmus paper to red “Blue to Red A-CID”

Some Properties of Bases Produce OH- ions in water Taste bitter, chalky Are electrolytes Feel soapy, slippery React with acids to form salts and water pH greater than 7 Turns red litmus paper to blue “Basic Blue”

Acid Nomenclature Review No Oxygen w/Oxygen An easy way to remember which goes with which… “In the cafeteria, you ATE something ICky”

Acid/Base definitions Definition 1: Arrhenius Arrhenius acid is a substance that produces H+ (H3O+) in water Arrhenius base is a substance that produces OH- in water

Acid/Base Definitions Definition #2: Brønsted – Lowry Acids – proton donor Bases – proton acceptor A “proton” is really just a hydrogen atom that has lost it’s electron!

A Brønsted-Lowry acid is a proton donor A Brønsted-Lowry base is a proton acceptor conjugate acid conjugate base base acid

ACID-BASE THEORIES The Brønsted definition means NH3 is a BASE in water — and water is itself an ACID

Conjugate Pairs

Learning Check! Label the acid, base, conjugate acid, and conjugate base in each reaction: HCl + OH-    Cl- + H2O Acid Base Conj.Base Conj.Acid H2O + H2SO4    HSO4- + H3O+ Conj.Base Conj.Acid Base Acid

Acids & Base Definitions Definition #3 – Lewis Lewis acid - a substance that accepts an electron pair Lewis base - a substance that donates an electron pair

Lewis Acids & Bases Formation of hydronium ion is also an excellent example. Electron pair of the new O-H bond originates on the Lewis base.

Lewis Acid/Base Reaction

The pH scale is a way of expressing the strength of acids and bases The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H+ (or OH-) ion. Under 7 = acid 7 = neutral Over 7 = base

(Remember that the [ ] mean Molarity) Calculating the pH pH = - log [H+] (Remember that the [ ] mean Molarity) Example: If [H+] = 1 X 10-10 pH = - log 1 X 10-10 pH = - (- 10) pH = 10 Example: If [H+] = 1.8 X 10-5 pH = - log 1.8 X 10-5 pH = - (- 4.74) pH = 4.74

Try These! pH = - log [H+] pH = - log 0.15 pH = - (- 0.82) pH = 0.82 pH = - log 3 X 10-7 pH = - (- 6.52) pH = 6.52 Find the pH of these: A 0.15 M solution of Hydrochloric acid 2) A 3.00 X 10-7 M solution of Nitric acid

pH calculations – Solving for H+ If the pH of Coke is 3.12, [H+] = ??? Because pH = - log [H+] then - pH = log [H+] Take antilog (10x) of both sides and get 10-pH = [H+] [H+] = 10-3.12 = 7.6 x 10-4 M *** to find antilog on your calculator, look for “Shift” or “2nd function” and then the log button

More About Water Equilibrium constant for water = Kw H2O can function as both an ACID and a BASE. In pure water there can be AUTOIONIZATION Equilibrium constant for water = Kw Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 oC

More About Water and so [H3O+] = [OH-] = 1.00 x 10-7 M Autoionization Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 oC In a neutral solution [H3O+] = [OH-] and so [H3O+] = [OH-] = 1.00 x 10-7 M

pOH Since acids and bases are opposites, pH and pOH are opposites! pOH does not really exist, but it is useful for changing bases to pH. pOH looks at the perspective of a base pOH = - log [OH-] Since pH and pOH are on opposite ends, pH + pOH = 14

pH [H+] [OH-] pOH

[H3O+], [OH-] and pH What is the pH of the 0.0010 M NaOH solution? [OH-] = 0.0010 (or 1.0 X 10-3 M) pOH = - log 0.0010 pOH = 3 pH = 14 – 3 = 11 OR Kw = [H3O+] [OH-] [H3O+] = 1.0 x 10-11 M pH = - log (1.0 x 10-11) = 11.00

What is the pH of a 2 x 10-3 M HNO3 solution? HNO3 is a strong acid – 100% dissociation. Start 0.002 M 0.0 M 0.0 M HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq) End 0.0 M 0.002 M 0.002 M pH = -log [H+] = -log [H3O+] = -log(0.002) = 2.7 What is the pH of a 1.8 x 10-2 M Ba(OH)2 solution? Ba(OH)2 is a strong base – 100% dissociation. Start 0.018 M 0.0 M 0.0 M Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq) End 0.0 M 0.018 M 0.036 M pH = 14.00 – pOH = 14.00 + log(0.036) = 12.56

Strong and Weak Acids/Bases The strength of an acid (or base) is determined by the amount of IONIZATION. HNO3, HCl, HBr, HI, H2SO4 and HClO4 are the strong acids.

Strong and Weak Acids/Bases Generally divide acids and bases into STRONG or WEAK ones. STRONG ACID: HNO3 (aq) + H2O (l)  H3O+ (aq) + NO3- (aq) HNO3 is about 100% dissociated in water.

Strong and Weak Acids/Bases Weak acids are much less than 100% ionized in water. *One of the best known is acetic acid = CH3CO2H

Strong and Weak Acids/Bases Strong Base: 100% dissociated in water. NaOH (aq)  Na+ (aq) + OH- (aq) Other common strong bases include KOH and Ca(OH)2. CaO (lime) + H2O --> Ca(OH)2 (slaked lime) CaO Strong bases are the group I hydroxides Calcium, strontium, and barium hydroxides are strong, but only soluble in water to 0.01 M

Strong and Weak Acids/Bases Weak base: less than 100% ionized in water One of the best known weak bases is ammonia NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq)

Weak Bases

Equilibria Involving Weak Acids and Bases Consider acetic acid, HC2H3O2 (HOAc) HC2H3O2 + H2O ↔ H3O+ + C2H3O2 - Acid Conj. base (K is designated Ka for ACID) K gives the ratio of ions (split up) to molecules (don’t split up)

Ionization Constants for Acids/Bases Conjugate Bases Increase strength Increase strength

Equilibrium Constants for Weak Acids Weak acid has Ka < 1 Leads to small [H3O+] and a pH of 2 - 7

Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. Step 1. Define equilibrium concs. in ICE table. [HOAc] [H3O+] [OAc-] initial change equilib 1.00 0 0 -x +x +x 1.00-x x x

Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. Step 2. Write Ka expression This is a quadratic. Solve using quadratic formula. or you can make an approximation if x is very small! (Rule of thumb: 10-5 or smaller is ok)

Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. Step 3. Solve Ka expression First assume x is very small because Ka is so small. Now we can more easily solve this approximate expression.

Equilibria Involving A Weak Acid You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH. Step 3. Solve Ka approximate expression x = [H3O+] = [OAc-] = 4.2 x 10-3 M pH = - log [H3O+] = -log (4.2 x 10-3) = 2.37

Equilibria Involving A Weak Acid Calculate the pH of a 0.0010 M solution of formic acid, HCO2H. HCO2H + H2O ↔ HCO2- + H3O+ Ka = 1.8 x 10-4 Approximate solution [H3O+] = 4.2 x 10-4 M, pH = 3.37 Exact Solution [H3O+] = [HCO2-] = 3.4 x 10-4 M [HCO2H] = 0.0010 - 3.4 x 10-4 = 0.0007 M pH = 3.47

Equilibrium Constants for Weak Bases Weak base has Kb < 1 Leads to small [OH-] and a pH of 12 - 7

Relation of Ka, Kb, [H3O+] and pH

Equilibria Involving A Weak Base You have 0.010 M NH3. Calc. the pH. NH3 + H2O ↔ NH4+ + OH- Kb = 1.8 x 10-5 Step 1. Define equilibrium concs. in ICE table [NH3] [NH4+] [OH-] initial change equilib 0.010 0 0 -x +x +x 0.010 - x x x

Equilibria Involving A Weak Base You have 0.010 M NH3. Calc. the pH. NH3 + H2O  NH4+ + OH- Kb = 1.8 x 10-5 Step 2. Solve the equilibrium expression Assume x is small, so x = [OH-] = [NH4+] = 4.2 x 10-4 M and [NH3] = 0.010 - 4.2 x 10-4 ≈ 0.010 M The approximation is valid !

Equilibria Involving A Weak Base You have 0.010 M NH3. Calc. the pH. NH3 + H2O  NH4+ + OH- Kb = 1.8 x 10-5 Step 3. Calculate pH [OH-] = 4.2 x 10-4 M so pOH = - log [OH-] = 3.37 Because pH + pOH = 14, pH = 10.63

Types of Acid/Base Reactions: Summary

Weak Bases are weak electrolytes F- (aq) + H2O (l) OH- (aq) + HF (aq) NO2- (aq) + H2O (l) OH- (aq) + HNO2 (aq) Conjugate acid-base pairs: The conjugate base of a strong acid has no measurable strength. H3O+ is the strongest acid that can exist in aqueous solution. The OH- ion is the strongest base that can exist in aqueous solution.

Strong Acid Weak Acid

For a monoprotic acid HA Ionized acid concentration at equilibrium Initial concentration of acid x 100% percent ionization = For a monoprotic acid HA Percent ionization = [H+] [HA]0 x 100% [HA]0 = initial concentration

Ionization Constants of Conjugate Acid-Base Pairs HA (aq) H+ (aq) + A- (aq) Ka A- (aq) + H2O (l) OH- (aq) + HA (aq) Kb H2O (l) H+ (aq) + OH- (aq) Kw KaKb = Kw Weak Acid and Its Conjugate Base Ka = Kw Kb Kb = Kw Ka

Molecular Structure and Acid Strength H X H+ + X- Bond strength Polarity The stronger the bond The weaker the acid HF << HCl < HBr < HI

Molecular Structure and Acid Strength Z O H O- + H+ d- d+ The O-H bond will be more polar and easier to break if: Z is very electronegative or Z is in a high oxidation state

Molecular Structure and Acid Strength 1. Oxoacids having different central atoms (Z) that are from the same group and that have the same oxidation number. Acid strength increases with increasing electronegativity of Z H O Cl O O • H O Br O O • Cl is more electronegative than Br HClO3 > HBrO3 15.9

Molecular Structure and Acid Strength 2. Oxoacids having the same central atom (Z) but different numbers of attached groups. Acid strength increases as the oxidation number of Z increases. HClO4 > HClO3 > HClO2 > HClO

Acid-Base Properties of Salts Neutral Solutions: Salts containing an alkali metal or alkaline earth metal ion (except Be2+) and the conjugate base of a strong acid (e.g. Cl-, Br-, and NO3-). NaCl (s) Na+ (aq) + Cl- (aq) H2O Basic Solutions: Salts derived from a strong base and a weak acid. NaCH3COO (s) Na+ (aq) + CH3COO- (aq) H2O CH3COO- (aq) + H2O (l) CH3COOH (aq) + OH- (aq)

Acid-Base Properties of Salts Acid Solutions: Salts derived from a strong acid and a weak base. NH4Cl (s) NH4+ (aq) + Cl- (aq) H2O NH4+ (aq) NH3 (aq) + H+ (aq) Salts with small, highly charged metal cations (e.g. Al3+, Cr3+, and Be2+) and the conjugate base of a strong acid. Al(H2O)6 (aq) Al(OH)(H2O)5 (aq) + H+ (aq) 3+ 2+

Acid-Base Properties of Salts Solutions in which both the cation and the anion hydrolyze: Kb for the anion > Ka for the cation, solution will be basic Kb for the anion < Ka for the cation, solution will be acidic Kb for the anion  Ka for the cation, solution will be neutral