Acids, Bases, and pH Acids, Bases, and pH Acids, Bases, and pH Acids, Bases, and pH.
Acids and Bases. What are acids? Examples? What are bases? Examples?
Transcript of Acids and Bases. What are acids? Examples? What are bases? Examples?
Acids and Bases
What are acids?
Examples?
What are bases?
Examples?
3 different definitions of acids/bases
• Arrhenius• Bronsted-Lowry• Lewis
Least broad
Most broad
Arrhenius
• Acids = a compound that increases the [H+] in aqueous solutions
• Bases = a compound that increases the
[OH-] in aqueous solutions
Ex of Arrhenius AcidH2O + HCl H3O+ + Cl-
Ex of Arrhenius Base
H2O + NH3 NH4+ + OH-
Limitations to Arrhenius’ Definition
• Aqueous = in water • Some acids and bases still act as acids or
bases even when they aren’t in water.
Bronsted-Lowry
• Acid = molecule or ion that is a PROTON DONOR
• Base = molecule or ion that is a PROTON ACCEPTOR
• Proton = H+
Example
NH3 + HCl NH4+ + Cl-
*Need to memorize that ammonia is NH3 and that it is a base
Conjugates
• Acid base reactions can go in reverse. • Each of the products can be classified as an
acid or base as well. • The species that started as an acid becomes
the conjugate base and vice versa
NH3 + HCl NH4+ + Cl-
• Which product is the conjugate acid? (can donate a H+)
• Which product is the conjugate base? (can accept a H+)
Amphoteric (Amphiprotic) Compounds
• Can act as either an acid or a base (donating or receiving an H+)
• WATER is a common example
• H2O + CH3COOH H3O+ + CH3COO-
• H2O + NH3 OH- + NH4+
Strengths of Acids/Bases
Strong Acids= easily lose protons (100%)
Weak acids= some protons are lost
Strong Bases= easily accept protons (100%)
Weak bases= some accept protons
If an acid is strong, the conjugate base is weak, and vice versa
Lewis
• Acid = an electron pair acceptor• Base = an electron pair donor• A + :B → A—B • H+ + :NH3 → NH4
+
• BF3 + F− → BF4−
Polyprotic acids
• Have multiple protons to lose• In excess base (in this case water)
• H3PO4 +H2O H2PO4 - +H3O+
• H2PO4 - + H2O HPO4 2- + H3O+
• HPO4 2- + H2O PO4 3- + H3O+
Polyprotic acids continued…
• Each time that a polyprotic loses an H+ it becomes harder to lose. Why?
• Therefore which acid in a polyprotic is the most acidic?
Prefixes: Di, Mono, Poly
• Monoprotic = only has one H+ to lose• Diprotic = has two H+ to lose (H2SO4)
• Polyprotic = has multiple (poly) H+ to lose
Homework
Chapter 16: #1,15,18,20,22,24,27,28
Autoionization of Water
Pure water self ionizes to a small extent
H2O H+ + OH-
H+ H3O+ (Hydronium ion) (Attaches onto a water molecule)
DYNAMIC equilibrium- no single molecule stays ionized for long
The amount it ionizes is very small.
In pure water [H3O+ ] = [OH-] = 1.00x10-7 M
K expression: (Kw stands for water ionization constant)
Kw = [H3O+ ] * [OH-]
K = [1.00x10-7 M ][1.00x10-7 M ] = 1.00x10-14 M (at 25 deg. Celsius)
Kw can be used to calculate hydronium ion or hydroxide ion concentrations at any time. Together their product is always 1.00x10-14M
If [H3O+] > [OH-] then the solution is acidic
If [OH-] > [H3O+] then the solution is basic
If they are equal (and therefore both 1.00 x 10-7 M) the solution is neutral
Example
Determine the hydronium ion concentration if a solution has a hydroxide concentraion of 0.00043M.
Is this an acidic, basic, or neutral solution?
pH
Hydronium power or potential
Negative Log scale of [H3O+] pure water has a pH of 7 because -log(1.00x10-7) = 7
Higher pH = lower concentration of Hydronium
Lower pH = higher concentration of H3O+
pOH is the same log scale, but for OH-
pOH + pH = 14
(because [OH-]*[H3O+]=1.00x10-14)
Helpful box
pH Convert using pH+pOH=14 pOH
Convert using: pH= -log[H3O+] Convert using: pOH= -log[OH-]
Or 10-pH = [H3O+] Or 10-pOH = [OH-]
[H3O+] Convert using Kw [OH-]
Examples
n What is the pH of a 0.040 M HCl solution?
n What is the pH of a 0.005M H2SO4 solution?
n What is the pH of a 0.008M Ca(OH)2 solution?
Homework
n 31,33,40,46,50