1 acids and bases operational definitions are based on observed properties. compounds can be...
DESCRIPTION
3 Properties of Bases Bases taste bitter; mustard and soap Bases cause weak organic acids (dyes) to change colour (red litmus paper to blue {BB} Basic Blue Acids destroy base properties - react with acids to form salts and water Bases are electrolytes {strong or weak} Feel soapy, slippery Bases are formed when the oxide of some metals dissolve in water (CaO(s) + H 2 O → Ca(OH) 2 (aq) {CaO is the base anhydride}TRANSCRIPT
1
Acids and Bases
Operational definitions are based on observed properties. Compounds can be Classified as acid or base by observing these sets of properties.
2
Properties of Acids Taste sour (acere – Latin for sour) (Lemons, vinegar)
Cause certain organic dyes to change colour (Turns blue litmus paper to red –
BAR)
Acid properties are destroyed by Bases (React with bases to form a salt and
water)
Acid solutions are Electrolytes (substance in solution that conduct an electric
current – Acids can be strong or weak electrolytes)
Acids react (corrode) with active metals (Group I and II as well as Zn and
Aluminum) (Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g))
Acids react with carbonates (CO32-) and hydrogen carbonates (HCO3
1-) to
produce carbon dioxide gas {2HCl(aq) + Na2CO3(s) → 2NaCl(aq) + H2O(l) + CO2(g)}
Certain nonmetal oxides will dissolve to produce acid solutions. (SO3(g) + H2O →
H2SO4(aq) (SO3(g) is the acid anhydride – without water)
3
Properties of Bases Bases taste bitter; mustard and soapBases taste bitter; mustard and soap
Bases cause weak organic acids (dyes) to change colour (red Bases cause weak organic acids (dyes) to change colour (red
litmus paper to blue {BB} Basic Bluelitmus paper to blue {BB} Basic Blue
Acids destroy base properties - react with acids to form salts Acids destroy base properties - react with acids to form salts
and water and water
Bases are electrolytes {strong or weak}Bases are electrolytes {strong or weak}
Feel soapy, slipperyFeel soapy, slippery
Bases are formed when the oxide of some metals dissolve in Bases are formed when the oxide of some metals dissolve in
water (CaO(s) + Hwater (CaO(s) + H22O O → Ca(OH)→ Ca(OH)22(aq)(aq) {CaO is the base anhydride} {CaO is the base anhydride}
4
Acid/Base definitions• Definition #1: Arrhenius (traditional)
– Acids are compounds with ionizable hydrogen– produce H+ ions (or hydronium ions H3O+) in solution
– Bases are compounds that produce OH- ions in solution (problem: some bases don’t have hydroxide ions!)
The reaction between an acid and a base:H+(aq) + OH-(aq) → H2O (l)
5
Arrhenius acid is a substance that produces H+ (H3O+) in water.The HCl molecule is ionized. (ionization)
Arrhenius base is a substance that produces OH- in water. The ions are dissociated. (dissociation)
6
Some acids have more than one ionizable hydrogen
• H2SO4 → H+(aq) + HSO41-(aq
• HSO41- → H+(aq) + SO4
2-(aq) H2SO4 is diprotic
• H3PO4(aq) → H+(aq) + H2PO41-(aq)
• H2PO41-(aq) → H+(aq) + HPO4
2-(aq)
• HPO42-(aq) → H+(aq) + PO4
3-(aq)
Phosphoric acid is a triprotic acid.
7
Water self-ionization
H2O ↔ H+(aq) + OH-(aq) [H+] = [OH-] = 10-7M at SATP
Keq = [H+][ OH-] [H2O(l)]Kw = [H+][ OH-] = 10-7 x 10-7 (at 25ºC)Kw = 10-14 at SATP
8
H2O ↔ H+(aq) + OH-(aq) What happens to this equilibrium if HCl(g) dissolves in the water? HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq)
• Increasing Decreasing• H2O ↔ H3O+(aq) + OH-(aq)• [H+] > [OH-] = acidic
• What happens when sodium hydroxide dissolves? NaOH(s) + H2O → Na+(aq) + OH-(aq)
• Decreasing Increasing• H2O ↔ H3O+(aq) + OH-(aq)• [H+] < [OH-] = basic (alkaline solution)
If [HIf [H++] = 10] = 10-7-7 then [OH then [OH--] = 10] = 10-7-7 solution is neutral (SATP) solution is neutral (SATP)
9
pH and logs• [H+] is important in the study of acid-base
chemistry. pH is the widely used scale to show [H+].
• pH = -log[H+] or pH = 1 . log[H+][H+] = 10 – pH (the antilog)A logarithm is the power to which ten must
be raised to get a number.log1000 = log(103) = 3
10
pH calculations• For a neutral solution • pH = -log[H+]• pH = -log [10-7]• pH = - [-7]• pH = 7 at SATP
• Example:• [H+] = 5 x 10-3
• pH = -log [5 x 10-3]• pH = -log [0.005]• pH = - (-2.3) = 2.3
11
pH and pOH• pOH = - log [OH-] or [OH-] = 10 - pOH
• Kw = [H+] x [OH-] = 1 x 10-14 (at 25ºC)• pKw= pH + pOH• 14= pH + pOH
• Example:• If pH = (2.3) what is the [OH-]?• pH + pOH = 14• pOH = 14 – pH• pOH = 14 – 2.3• pOH = 11.7• pOH = -log [OH-]• [OH-] = inverse log -11.7 or (10 - 11.7)• [OH-] = 2.0 x 10-12
12
[H[H33OO++], [OH], [OH--] and pH] and pH• What is the pH of the 0.0010 M NaOH What is the pH of the 0.0010 M NaOH
solution? solution? • [OH[OH--] = 0.0010 (or 1.0 X 10] = 0.0010 (or 1.0 X 10-3-3 M) M)• pOH = - log 0.0010pOH = - log 0.0010• pOH = 3pOH = 3• pH + pOH = 14pH + pOH = 14• pH = 14 – 3 = 11pH = 14 – 3 = 11• OR Kw = [HOR Kw = [H33OO++] [OH] [OH--]]• 1.0 x101.0 x10-14-14 = [H = [H3OO++] x 1.0 X 10] x 1.0 X 10-3-3 • [H[H3OO++] = 1.0 x 10] = 1.0 x 10-11-11 M M• pH = - log (1.0 x 10pH = - log (1.0 x 10-11-11) = 11.00) = 11.00
13
Problem 1: The pH of rainwater collected in a certain region of the northeastern New Brunswick on a particular day was 4.82. What is the H+ ion concentration of the rainwater?
Problem 2: The OH- ion concentration of a blood sample is 2.5 x 10-7M. What is the pH of the blood?
Problem 3: A chemist dilutes concentrated hydrochloric acid to make two solutions: (a) 3.0 M and (b) 0.0024 M. Calculate the [H3O+], pH, [OH-], and pOH of the two solutions at 25°C.
[H+] = 1.51 x 10-5
pOH = 6.6 pH = 7.4
Problem 4: What is the [H3O+], [OH-], and pOH of a solution with pH = 3.67? Is this an acid, base, or neutral?
Problem 5: Problem #4 with pH = 8.05?
a) [H3O+] = [3.0], pH = - 0.48, pOH = 14.48, [OH-] = 3.3 x 10-15
b) [H3O+] = [2.4x10-3], pH = 2.62, pOH = 11.38, [OH-] = 4.2 x 10-12
[H3O+] = 2.14 x10-4, pOH = 10.33, [OH-] = 4.68x 10-11 It is an acid.
[H3O+] = 8.92 x10-9, pOH = 5.95, [OH-] = 1.12x 10-6 It is an acid.
14
Acid/Base Definitions
• Definition #2: Brønsted – Lowry
– Acids – proton donor A “proton” is a hydrogen ion (the atom lost it’s electron)
– Bases – proton acceptor (accepts a hydrogen ion) No longer needs to contain the OH- ion
15
A Brønsted-Lowry acid is a proton donorA Brønsted-Lowry base is a proton acceptor
acid conjugate base
base conjugate acid
BaseAcidAcidBaseNH4
+ + OH-NH3 + H2O
16
The Bronsted-Lowry conceptThe Bronsted-Lowry concept
• Acids and bases are identified based on whether they donate or accept H+.
• “Conjugate” acids and bases are found on the products side of the equation. A conjugate base is the same as the starting acid minus H+.
+Cl HH
HO
+H
HH O Cl+
acid base conjugate acid conjugate base
conjugate acid-base pairs
17
Practice problemsPractice problemsIdentify the acid, base, conjugate acid, conjugate base, and conjugate acid-base pairs:
acid base conjugate acidconjugate baseCH3OOH(aq) + H2O(l) CH3COO–(aq) + H3O+(aq)
conjugate acid-base pairs
acidbase conjugate acidconjugate baseOH
–(aq) + HCO3–(aq) CO3
2–(aq) + H2O(l)
conjugate acid-base pairs
18
Base Conjugate acid \ \NH3(g) + H2O(l) ↔ NH4
+(aq) + OH-(aq) / / Acid Conjugate Base
HCl(aq) + H2O(l) ↔ H3O+(aq) + Cl-(aq)Acid Base Conjugate Conjugate
Acid Base
The water has acted as both an acid and a base, depending on what it is mixed with. Substances that can act as both an acid and a base are amphoteric (also called amphiproteric).
19
Strong acid and base :HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq)B(aq) + H2O(l) ↔ BH+(aq) + OH-(aq)At equilibrium the ionic form is favored
Weak acid and base :HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq)
B(aq) + H2O(l) ↔ BH+(aq) + OH-(aq)
At equilibrium the molecular form is favored
20
CH3COOH(aq) + H2O(l) ↔ H+(aq) + CH3COO-(aq)
Keq= [H+] [CH3COO-] . [CH3COOH] [H2O(l)][H2O] is a constant, so collect the constants(Keq)[H2O(l)] = [H+] [CH3COO-]
[CH3COOH] (Keq)[H2O] is represented Ka(ionization constant for an acid) Ka = [H+] [CH3COO-] = 1.8 x 10-5
[CH3COOH]Ka < 1 weak acidGeneral Formula for the ionization constant of a weak acid.
21