Unit 12 Acids and Bases
The stronger the acid...
the weaker its conjugate base
Each reaction has...
two acids and two bases
Weak acids are...
weaker than H₃O⁺
Why do strong acids have higher Ka's?
Favors products: More dissociates, bigger Ka (higher ability to give H⁺)
Kb
base dissociation constant
Amphoteric
can act as both an acid and a base
What happens when the structure of indicator changes when exposed to different pH?
changes color
Acid-base reactions will always...
drive towards the weak side (Right: Kc>1, Left: Kc<1)
Conjugate base of weak acid...
is even weaker
Stronger acid =
larger Ka
A conjugate base has...
one fewer H and one more negative charge than the acid
A conjugate acid has...
one more H and one less negative charge than the base
pKw
pH +pOH =14.00 (only at 25°C)
Water turns all acids ____________ than hydronium into hydronium
stronger
Water turns all bases ____________ than hydroxide into hydroxide
stronger
Acid-base reactions proceed from...
stronger acid and base to weaker acid and base
Strong acids are...
stronger than H₃O⁺
Weak bases
-Ammonia (NH₃) -Organic amines (RNH₂, R₂NH, R₃H) -Anything basically with a lone pair Examples: CH₃CH₂NH₂, (CH₃)₂NH, (CH₃)₃N
Polyprotic acids
-Have more than one acidic proton -The 1st dissociation is all that really matters/significant -In every case for a polyprotic acid: Ka₁>Ka₂>Ka₃
Strong acids
-Hydrohalic acids (except HF) -Oxyacids where number of O atoms exceeds the number of H atoms by two or more Examples: H₂SO₄ is strong (HSO₄⁻ is weak CB); H₂CO₃ is weak (HCO₃⁻ is even weaken CB)
Strong bases
-M₂O or MOH, where M = Any Group IA metal (soluble) -MO or M(OH)₂, where M = Ca, Sr, Ba (Lower IIA metals)
Weak acids
-Oxyacids that don't fit the strong acid O to H rule (anything else) -Organic acids Examples: HF, HCN, H₂S, R-COOH
(using [H₃O⁺]) pH =
-log[H₃O⁺]
(using [OH⁻]) pOH =
-log[OH⁻]
Categories of Lewis Acids
1. Electron deficient central atom molecules (could accept a pair of e⁻), contains B and Al 2. Molecules with polar double bonds, double bond becomes single bond, new bonds forms. Examples are SO₂, H₂O, CO₂ 3. Metal cations, complexation reactions, hydrated metal ions. Example is Al³⁺
Kw @ 25°C
1.0 x 10⁻¹⁴
(using pH) [H₃O⁺] =
10^-pH
(using pOH) [OH⁻] =
10^-pOH
Weak base dissociation equation
B(aq) + H₂O(l)⇌BH⁺(aq) +OH⁻(aq)
Will an aqueous solution of (NH₄)₂CO₃ at 25°C be acidic or basic?
BASIC: Compare the Ka of NH₄⁺ and Kb of CO₃⁻². Take Kw/Kb of NH₃ to find Ka of NH₄⁺. Take Kw/Ka of HCO₃⁻ to find Kb of CO₃⁻². Kb of CO₃⁻² is much stronger than Ka of NH₄⁺, so the solution will be basic.
Acid of CH₃OO⁻ (conjugate base)
CH₃COOH
Lewis acids
Electron pair acceptors; it broadens our definition of acids
Lewis bases
Electron pair donors; lewis bases and B-L bases are the same
Acid/base strength chart
FOR ACIDS/LEFT SIDE: Increasing acid strength/Ka from bottom to top; strongest acids at top/weakest acids at bottom FOR BASES/RIGHT SIDE: Increasing base strength/Kb from top to bottom; strongest bases at bottom/weakest bases at top Those that appear twice are amphoteric
Strong acid dissociation equation
HA(aq) + H₂O(l)→H₃O⁺(aq) + A⁻(aq)
Weak acid dissociation equation
HA(aq) + H₂O(l)⇌H₃O⁺(aq) + A⁻(aq)
Comparing the strength of HCl and HBr by dissolving them in a solvent that is a weaker base than H₂O
HCl(ac) + HAc⇌H₂Ac⁺(ac) + Cl⁻(ac) Kc₁ HBr(ac) + HAc⇌H₂Ac⁺(ac) + Cl⁻(ac) Kc₂ Kc₂ > Kc₁, so HBr is the stronger acid
Conjugate base of H₂S
HS⁻
Diprotic
Has more than one H⁺ ion to lose
Base of H₃O⁺ (conjugate acid)
H₂O
A scientist finds the pH of a 0.12 M HPAc solution 2.60. Calculate the Ka/percent dissociation of HPAc..l
Ka = 5.2 x 10⁻⁵ % Dissociation = 2.1% HOW TO: Write the equation (HPAc + H₂0⇌H₃O⁺ + PAc⁻). Set up an ICE table. Use 10^-pH to find [H₃O⁺]. Plug in that number for x and square it. Divide it by 0.12 to find the Ka. Take the x value, divide it by 0.12, and multiply by 100 to find the % dissociation.
The value of Ka for HCN is 4.0 x 10⁻¹⁰. What is the Kb for the CN⁻ ion at 25°C?
Kb = 2.5 x 10⁻⁵ HOW TO: Since Ka x Kb = Kw, take Kw/Ka to find the Kb
Ka x Kb =
Kw
As temperature increases...
Kw increases
Identify the Lewis acid and base: Co³⁺ + 6NH₃→Co(NH₃)₆³⁺
Lewis acid: Co³⁺ (highly charged metal ion) Lewis base: 6NH₃ (can donate lone pair)
Identify the Lewis acid and base: SO₃ + H₂O→H₂SO₄
Lewis acid: SO₃ (nonmetal oxide) Lewis base: H₂O
Identify the Lewis acid and base: OH⁻ + Al(OH)₃→Al(OH)₄⁻
Lewis base: OH⁻ (ligand, has to have a lone pair) Lewis acid: Al(OH)₃
Conjugate acid if NH₃
NH₄⁺
Effects of salts of strong acids and strong bases in water
No change in pH
Dissolve NaCl in water
No pH change: Na⁺ comes from a strong base (CA of NaOH), so it is a negligible acid. Cl⁻ comes from a strong acid (CB of HCl), so it is a negligible base. Equation: HCl + NaOH→H₂O + NaCl (Na⁺/Cl⁻)
Left ot right periodic trend for acid strength
Normal way: Compare their CB's, and since they are very similar in size, the most electronegative one is the strongest College Board/Test: The more polar the H-X bond, the more acidic the compound *Increasing acid strength from right to left*
Top to bottom periodic trend for acid strength
Normal way: The bigger the CB, the more it is able to spread out the charge College Board/Test: The weaker/longer the H-X bond, the more acidic the compound *Increasing acid strength from top to bottom*
Oxyacid strength trend: A series of oxyacids, acidity increases with the number of oxygens (why?)
The more oxygens there are, the more stable/more resonance there is in it's CB
Oxyacid strength trend: -OH bonded to another atom Y (nonmetal)
The more electronegative Y is, the more acidic the acid
What can responsible for color changes in indicators?
Transition metal with ligands or conjugated organic compounds
Kb =
[HB][OH⁻]/[B⁻]
Percent dissociation formula
[H₃O⁺] (your x value) / [HA] X 100 (less than 5%)
Basic solutions
[H₃O⁺] < [OH⁻] (higher pH)
Propanoic acid, HPr, is an organic acid whose salts are used to slow mold growth in foods. What is the [H₃O⁺] and pH of 0.10 M HPr if the Ka = 1.3 x 10⁻⁵?
[H₃O⁺] = 1.1 x 10⁻³ M pH = 2.96 HOW TO: Write the equation (HPr + H₂O⇌H₃O⁺ + Pr⁻). Set up an ICE table. Set the Ka value equal to x²/0.10 and find x/[H₃O⁺]. Do -log[H₃O⁺] to find the pH.
Calculate [H₃O⁺] in a solution at 25°C whose [OH⁻] is 6.7 x 10⁻² M. Is the solution acidic, basic, or neutral?
[H₃O⁺] = 1.5 x 10⁻¹³ Basic: pH = 12.83 HOW TO: Use Kw = [H₃O⁺][OH⁻] to find [H₃O⁺]. Then do -log[H₃O⁺] to find the pH.
Neutral solutions
[H₃O⁺] = [OH⁻]
Since the ratio of hydronium and hydroxide is 1:1...
[H₃O⁺] = [OH⁻] (usually you'll take square root/x²)
Acidic solutions
[H₃O⁺] > [OH⁻] (lower pH)
Ka =
[H₃O⁺][A⁻]/[HA]
Kw =
[H₃O⁺][OH⁻]
A base must contain...
a lone pair of electrons to bind the H⁺ ion
a.) Calculate the pH of a solution that contains 1.00 M HCN (Ka = 6.2 x 10⁻¹⁰) and 5.00 M HNO₂ (Ka = 4.0 x 10⁻⁴) b.) Determine the concentration of cyanide ion at equilibrium (common ion effect) c.) What percent of nitrous acid dissociated (ionized)?
a.) pH = 1.35 HOW TO: Ignore the HCN since it is weaker than HNO₂. Write the equation (HNO₂ + H₂O⇌H₃O⁺ + NO₂⁻). Set up an ICE table. Set the Ka equal to x²/5.00 to find [H₃O⁺]. Take -log[H₃O⁺] to find the pH. b.) [CN⁻] = 1.4 x 10⁻⁸ M HOW TO: Write the equation (HCN + H₂O⇌H₃O⁺ + CN⁻). Set up an ICE table using 1.00 M for HCN and 0.045 M for H₃O⁺. Set the Ka value equal to 0.045x. X is your [CN⁻]. c.) 0.9% dissociation HOW TO: Take the [H₃O⁺]/[HNO₂] x 100 to get the percent dissociation.
Ka
acid dissociation constant
Dimethylamine, (CH₃)₂NH, has a Kb of 5.9 x 10⁻⁴. What is the pH of 1.5 M (CH₃)₂NH at 25°C (and percent dissociation)?
pH = 12.47 % Dissociation = 2% HOW TO: Write the equation ((CH₃)₂NH + H₂O⇌OH⁻ + (CH₃)₂NH₂⁺). Set up an ICE table. Set the Ka value equal to x²/1.5 to find the [OH⁻]. Find the pOH by taking -log[OH⁻]. Do 14 - pOH to find the pH. Take the x value/1.5 M x 100 to get the percent dissociation.
Ascorbic acid, H₂Asc, is a diprotic acid (Ka₁ = 1.0 x 10⁻⁵, Ka₂ = 5.0 x 10⁻¹²). Calculate the pH of 0.050 M solution of ascorbic acid (and percent dissociation to check).
pH = 3.15 % Dissociation = 1.4% HOW TO: Write the equation (H₂Asc + H₂O⇌H₃O⁺ + HAsc⁻). Set up an ICE table. Set the Ka value of 1st dissociation equal to x²/0.050 to find [H₃O⁺]. Take the -log[H₃O⁺] to find the pH. Take the x value/0.050 M x 100 to get the percent dissociation.
Calculate the pH of a 0.20 M solution of NH₄NO₃ at 25°C when the Kb of NH₃ = 1.8 x 10⁻5 (and percent dissociation to check)
pH = 4.95 % Dissociation = 0.0055% HOW TO: Determine that NH₄⁺ is the CA of a weak base so it has moderate acidity. Calculate the Ka from Kw/Kb. Write the equation (NH₄⁺ + H₂O⇌H₃O⁺ + NH₃). Set up an ICE table. Set the Ka value equal to x²/0.20 to find [H₃O⁺]. Take the -log[H₃O⁺] to find the pH. Take the x value/0.20 x 100 to get the percent dissociation.
Effect of salts of weak acids and weak bases in water
pH of solution depends on the relative acid and base strength of the ions (Ka vs Kb-acid more acidic vs base more basic) Example: HC₂H₃O₂ + NH₃→NH₄C₂H₃O₂; NH₄⁺ is CA of NH₃/weak base so it has moderate acidity, C₂H₃O₂⁻ is CB of HC₂H₃O₂/weak acid so it is moderately basic.
Effect of salts of strong acid and weak base in water
pH will decrease (more acidic)
Dissolve NH₄Cl in water
pH will decrease: Cl⁻ comes from a strong acid (CB of HCl), so it is a negligible base. NH₄⁺ comes from a weak base (CA of NH₃), so it has moderate acidity. Equation: NH₄⁺(aq) + H₂O(l)⇌NH₃(aq) + H₃O⁺(aq)
Effect of salts of weak acids and strong bases in water
pH will increase (more basic)
Dissolve NaC₂H₃O₂ in water
pH will increase: Na⁺ comes from a strong base (CA of NaOH), so it is a negligible acid. C₂H₃O₂⁻ comes from a weak acid (CB of HC₂H₃O₂), so it is moderately basic. Equation: C₂H₃O₂⁻(aq) + H₂O(l)⇌HC₂H₃O₂(aq) + OH⁻(aq)
pOH for solution of NaOH with pH of 9.52? [H₃O⁺]? [OH⁻]?
pOH = 4.48 [H₃O⁺] = 3.0 x 10⁻¹⁰ M [OH⁻] = 3.3 x 10⁻⁵ M HOW TO: Subtract 9.52 from 14 to find the pOH. Do 10^-pOH to find [OH⁻]. Do 10^-pH to find [H₃O⁺].
Bronsted-Lowry base
proton acceptor (H+)
Bronsted-Lowry acid
proton donor (H+)
Weaker acid =
smaller Ka