Acids and Bases Chapter 15

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”

Practice: H2S HNO2 H2CO3 HC2H3O2 HF H2O Hydrosulfuric acid X = sulfide Hydrosulfuric acid Nitrous acid X = nitrite Carbonic acid X = carbonate Acetic acid X = acetate Hydrofluoric acid X = fluoride Water is NOT an Acid pH = 7 Dihydrogen monoxide X = oxide

Naming Bases A base is an ionic compound that produces hydroxide ions when dissolved in water. Bases are named in the same way as other ionic compounds The name of the cation, followed by the name of the anion Anion will always be hydroxide (OH1-) Exception: NH3 - ammonia

Practice: LiOH Al(OH)3 Fe(OH)3 Mg(OH)2 Pb(OH)2 HOH Lithium hydroxide Aluminum hydroxide Iron (III) hydroxide Magnesium hydroxide Lead (II) hydroxide Water is NOT a Base pH = 7 Dihydrogen monoxide

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 4.3

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

Conjugate Acid-Base Pairs In the reaction of NH3 and H2O, one conjugate acid-base pair is NH3/NH4+ the other conjugate acid-base is H2O/H3O+.

Learning Check A. Write the conjugate base of the following. 1. HBr 2. H2S 3. H2CO3 B. Write the conjugate acid of the following. 1. NO2- 2. NH3 3. OH-

Solution A. Remove H+ to write the conjugate base. 1. HBr Br- 2. H2S HS- 3. H2CO3 HCO3- B. Add H+ to write the conjugate acid. 1. NO2- HNO2 2. NH3 NH4+ 3. OH- H2O

Learning Check 1. HNO2, NO2− 2. H2CO3, CO32− 3. HCl, ClO4− 4. HS−, H2S Identify the sets that contain acid-base conjugate pairs. 1. HNO2, NO2− 2. H2CO3, CO32− 3. HCl, ClO4− 4. HS−, H2S 5. NH3, NH4+

Solution 1. HNO2, NO2− 4. HS−, H2S 5. NH3, NH4+ Identify the sets that contain acid-base conjugate pairs. 1. HNO2, NO2− 4. HS−, H2S 5. NH3, NH4+

Learning Check A. The conjugate base of HCO3− is 1. CO32− 2. HCO3− 3. H2CO3 B. The conjugate acid of HCO3- is 1. CO32− 2. HCO3− 3. H2CO3 C. The conjugate base of H2O is 1. OH− 2. H2O 3. H3O+ D. The conjugate acid of H2O is

Solution A. The conjugate base of HCO3 − is 1. CO32− B. The conjugate acid of HCO3− is 3. H2CO3 C. The conjugate base of H2O is 1. OH− D. The conjugate acid of H2O is 3. H3O+

3 Definitions of Acids & Bases 1) Arrhenius Theory Acids: ionize to produce H+ ions in aqueous solution Monoprotic: HNO3, HCl, HC2H3O2 Diprotic: H2SO4, H2SO3 Triprotic: H3PO4 Bases: dissociate to produce OH- ions in aqueous solution NaOH, KOH, Ca(OH)2, Al(OH)3

Definitions (cont’d) 2) Brønsted-Lowry Theory Hydrogen ion (H+) = a proton Acids: proton donors – ex. HCl Bases: proton acceptors – ex. NH3

Conjugate Acids and Bases Show the direction of H+ transfer. Label: Acid, Base, Conjugate Base, Conjugate Acid NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq) HCl (g) + H2O (l) Cl- (aq) + H3O+ (aq)

More Examples + WS Show the direction of H+ transfer. Label: Acid, Base, Conjugate Base, Conjugate Acid H2SO4 + OH- HSO41- + H2O HSO41- + H2O SO42- + H3O+

Conjugate Acid/Base Pairs WS

Definitions (cont’d) 3) Lewis Theory Acids: electron-pair acceptor Bases: electron-pair donor   HCl (g) + H2O (l) H3O+ (aq) + Cl- (aq)

pH Scale

pH Scale pH Scale: logarithmic scale in which [H3O+] is expressed as a number from 0 to 14. [OH-] = [H3O+] when pH = 7.

pH Scale Acidic Neutral Basic 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 100 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 10-10 10-11 10-12 10-13 10-14 [H3O+] 10-14 10-13 10-12 10-11 10-10 10-9 10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1 100 [OH-]

What does the pH scale mean?  Neutral Solution: [H3O+] = [OH-] A pH 11 solution has ____ times more hydroxide ions than a pH 10 solution. A pH 11 solution has _____ times more hydroxide ions than a pH 9 solution. A pH 11 solution has _____ the number of hydronium ions compared to a pH 10 solution.   10, 100, 1/10

pH and pOH table WS

pH – Examples (no calculator) What is the pH if [HCl] = 1 x 10-4 M? What is the [H+] if the pH = 9? What is the pH if [NaOH] = 1 x 10-2 M What is the concentration of [OH-] if the pOH is 3? What is the concentration of [H+] if the pOH is 10?

pH calculations – easy practice What is the pH of the solution? [H3O+] = 1 x 10-4 M [H+] = 1 x 10-10 M [HCl] = 1 x 10-2 M What is the concentration of H3O+ if the pH is 5? What is the concentration of H+ if the pH is 11?

Warm Up for Wednesday, May 22-Grab goggles! What is the pOH of each solution? [OH-] = 1 x 10-4 M [NaOH] = 1 x 10-10 M What is the pH of a solution if the pOH is 4? What is the pH of each solution? [OH-] = 1 x 10-8 M [KOH] = 1 x 10-3 M What is the [H3O+] in a solution if [OH-] = 1 x 10-3 M? What is the [OH-] in a solution if [H3O+] = 1x 10-5 M? What is the pH of [HCl] = 0.001 M?

pH equations pH= - log [H3O+] pOH = - log [OH-] pH + pOH = 14 [H+] x [OH-] = 1 x 10-14

Warm Up for the test! Find the pH of the following solutions. Is the solution acidic or basic? 0.01 M HCl 2.6 x 10-12 M Mg(OH)2 What would be considered a nonelectrolyte?. Find the concentration of hydroxide ions if the pH is 5.6. What would make a chemical reaction react quickly? Name or write the formula for the following: Hydrochloric acid Strontium hydroxide Nitric acid H2SO4 Ca(OH)2

On your scantron: Name: Date: 5/24/13 Period: Subject: Solutions/Acids/Bases When you are finished, work on the final exam study guide!

Strong Acids and Bases Strong Acids: completely ionize in water Ex. HCl, HBr, HI, HNO3, H2SO4, HClO3 Strong Bases: completely dissociate into ions in water Ex. NaOH, LiOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

Weak Acids and Bases Weak Acids: only some molecules ionize in water Ex: acetic acid (less than 0.5% of molecules ionize) Weak Bases: do not completely dissociate into ions in water Ex: ammonia (only 0.5% of molecules dissociate)

Concentrated vs. Strong “Concentrated” – refers to the amount dissolved in solution. “Strong” – refers to the fraction of molecules that ionize.   For example, if you put a lot of ammonia into a little water, you will create a highly concentrated solution. However, since only 0.5% of ammonia molecules ionize in water, this basic solution will not be very strong.

Neutralization Reaction ACID + BASE  SALT + WATER Salt: ionic compound formed from the negative part of the acid and the positive part of the base. Example:   2HCl + Mg(OH)2 MgCl2 + 2H2O What type of reaction is this? Synthesis, Decomposition, Single Replacement, Double Replacement, or Combustion?

Warm Up: Complete and balance the equations(produces salt and water) HCl + NaOH  HC2H3O2 + Ca(OH)2  HBr + Al(OH)3 

Acid-Base Titration Uses a neutralization reaction to determine the concentration of an acid or base. Standard Solution: the reactant that has a known molarity Endpoint: the point at which the unknown has been neutralized.

Titration Examples Example #1) 8.0 mL of 0.100 M NaOH is used to neutralize 20.0 mL of HCl. What is the molarity of HCl? NaOH + HCl  NaCl + H2O

Titration Examples (cont’d) Example #2) A 0.1 M Mg(OH)2 solution was used to titrate an HBr solution of unknown concentration. At the endpoint, 21.0 mL of Mg(OH)2 solution had neutralized 10.0 mL of HBr. What is the molarity of the HBr solution?

Titration Practice What is the molarity of an Al(OH)3 solution if 30.0 mL of the solution is neutralized by 26.4 mL of a 0.25 M HBr solution? A 0.3 M Ca(OH)2 solution was used to titrate an HCl solution of unknown concentration. At the endpoint, 35.0 mL of Ca(OH)2 solution had neutralized 10.0 mL of HCl. What is the molarity of the HCl solution? When 34.2 mL of a 1.02 M NaOH solution is added from a buret to 25.00 mL of a phosphoric acid solution that contains phenolphthalein, the solution changes from colorless to pink.  What is the molarity of the phosphoric acid?

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 15.4

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. 15.4

15.4

Strong Acid Weak Acid 15.4

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 15.5

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 15.7

Molecular Structure and Acid Strength H X H+ + X- Bond strength Polarity The stronger the bond The weaker the acid HF << HCl < HBr < HI 15.9

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 15.9

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 15.9

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) 15.10

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+ 15.10

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 15.10