1. Chemistry – April 21, 2017 P3 Challenge –
Consider: 2 SO2 (g) + 1 O2 (g)  2 SO3 (g) (Exothermic) at equilibrium.
Determine which way the equilibrium will shift when the following disturbances are applied.
A) SO3 is removed.
B) Temperature is increased .
C) Volume is increased.
Get out LeChatelier’s Principle
worksheet for HMW check

2. Chemistry – April 21, 2017 Today’s Objective –
Acid/Base equilibrium and pH
Assignment:
Acid/Base pH Worksheet
Agenda
Acid/base properties
Acid/base definitions
Acid/base concentrations
Dissociation and Ionization
Self-ionization of water
pH, pOH

3. Already about acids and bases (U5)
Acids react with bases to form water and a salt.
Type of Double Replacement reaction called a neutralization.
Forming water is a driving force
Acidic compounds start with H and have 3 naming rules.
There are 6 strong acids: HCl, HBr, HI, H2SO4, HNO3 and HClO4
All other acids are weak.
Strong bases are Group 1 and 2 hydroxides.

4. Acid and base properties
Acids properties:
Sour taste
Dissolves metals
Turns litmus paper red
Turns phenolphthalein colorless
Base properties
Bitter taste
Slippery feel
Turns litmus paper blue
Turns phenolphthalein pink

5. Arrhenius Definitions
Proposed by Svante Arrhenius, a Swedish scientist, in 1850
Acid = produces H+ ion in solution
Base = produces OH- ion in solution
Can label any substance based on its chemical formula.
Ex: NaOH
Ex: H2SO4
Limited usefulness because it misses some substances that we can label as acids and bases based on their properties: NH3 is a base, NaHCO3 is a base

6. Brønsted-Lowry Definition
Proposed independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923
Acid = proton (H+) donor
Base = proton (H+) acceptor
Some molecules can both donate and accept protons (Ex: Water) and are called amphoteric.
Ex: NH3(aq) + H2O (l)  NH4+ (aq) + OH- (aq)
Ex: HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq)

7. Types of acids Monoprotic acids donate 1 proton
Ex: HCl
Diprotic acids can donate 2 protons
Ex: H2SO4
Triprotic acids can donate 3 protons
Ex: H3PO4

8. Conjugate acids and Conjugate bases
The reactions of acids and bases are often reversible creating an equilibrium
The “acid product” that forms from the addition of a proton to a base is a conjugate acid.
The “base product” that forms from the removal of a proton from an acid is a conjugate base.
Any reactant and its conjugate partner form a conjugate acid-base pair.
Ex: NH3(aq) + H2O (l)  NH4+ (aq) + OH- (aq)

9. Dissociation vs Ionization
When an ionic compound dissolves in water = dissociation.
The ions already exists, they just move apart.
Ex: NaOH a strong base added to water is a dissociation.
When a molecular compound splits apart and forms ions = ionization
Ions must be created.
Ex: HC2H3O2 a weak acid, added to water is a partial ionization.

10. Strong Acid Concentrations
6 Strong acids completely ionize in water and are strong electrolytes.
HCl, HBr, HI,
HNO3, HClO4, H2SO4
All other acids are weak, only partially ionize. Form an equilibrium.
HA (aq)  H+ (aq) + A- (aq)
For strong acids, determine ion concentrations to estimate [H+]
1.5 M HCl
3.5 M H2SO4

11. Strong Base Concentrations
Strong bases are the hydroxides of alkali and alkaline earth metals.
Strong bases are strong electrolytes and dissolve completely in water.
All other bases are weak and only partially dissolve.
Determine the [OH-] concentration for
0.45 M Ba(OH)2

12. Water Because water is amphoteric: can act as an acid or a base.
Water can react with itself, acting as both the acid and base:
H2O(l) + H2O(l)  H3O+ (aq) + OH- (aq)
This reaction is the self-ionization of water, and happens in very small amounts.
In pure water at 25◦C, [H+] = [OH-] = 1 x 10-7 M = M
Product of these two concentrations is called the ion product of water Kw = 1 x = [H+] [OH-]
The ion product of water is true for all acidic or basic solutions.
H3O+ = H+

13. Relative concentrations
[H+]
[OH-] Kw
1.E-01
1.E-13
1.E-14
1.E-02
1.E-12
1.E-03
1.E-11
1.E-04
1.E-10
1.E-05
1.E-09
1.E-06
1.E-08
1.E-07

14. pH pH = -log [H+] [H+] Opp of Exp pH 1.E-01 1 1.E-02 2 1.E-03 3 1.E-045
1.E-06 6
1.E-07 7
1.E-08 8
1.E-09 9
1.E-10 10
1.E-11 11
1.E-12 12
1.E-13 13
Easier to refer to [H+] based on the exponent.
More specifically the opposite of the exponent because they are negative.
We call this value pH.
Small pH values, have highest [H+]
Mathematically,
pH = -log [H+]

15. Acidic, Basic, or Neutral pH
pH is by far the most common way to describe the acidity/basicity of a solution
Always remember that pH is a log function: each increase by 1 is a factor of 10.
Ex: pH = 7 pH = pH = 8.2

16. pOH It is less common than pH, but you can similarly define pOH
pOH = -log [OH-]
Low pOH values are basic, high pOH values are acidic, 7 is neutral.
pH +pOH = 14
Ex: If pOH = 4, pH =
Ex: If pOH = 10, pH =
Often easiest, to convert pOH to pH when identifying acidity.

17. pH Calculations Four related quantities: pH, pOH, [H+] and [OH-]
pH = -log [H+] , pOH = -log [OH-], Kw = [H+] [OH-] = 1 x , pOH + pH = 14
Note: the opposite of pH = -log [H+] is [H+] = 10 –pH
Calculate the other three quantities for the given information. Then indicate if it is an acidic or basic solution.
[H+] = 4.0 x 10-3 M
[OH-] = 1.5 x 10-5 M
pH = 2.1
pOH = 3.5

18. Exit Slip - Homework
Exit Slip: An acidic solution has [H+] = 4.5 x 10-4 M. What are a) the pH b) the pOH and c) [OH-]
What’s Due? (Pending assignments to complete.)
Acid/Base pH worksheet
What’s Next? (How to prepare for the next day)
Read p