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