1. Net Ionic Equations HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
What really happens:
H+(aq) + OH-(aq)  H2O(l)
Sodium ion and chloride ion are “spectator ions”

2. Reactions involving weak bases
HCl(aq) + NH3(aq)  NH4+(aq) + Cl-(aq)
Net-Ionic Equation:

3. What are the products: CH3CO2H(aq) + NaOH(aq) 
1. CH3CO2H2+(aq) + NaO(aq)
2. CH3CO2-(aq) + H2O(l) + Na+(aq)
3. CH4(g) + CO2(g) + H2O(l)
Net-Ionic Equation:

4. What are the products: HCN(aq) + NH3(aq) 
1. NH4+(aq) + CN-(aq)
2. H2CN+(aq) + NH2-(aq)
3. C2N2(s) + 3 H2(g)
Net-ionic equation:

5. Solution Concentration: Molarity
Molarity = moles solute per liter of solution
0.30 mol NH3 dissolved in L Concentration =
Written like: [NH3] = 0.60 M
If you dissolve 0.32 mol NH3 to make 750 mL of solution, what is [NH3]?

6. pH: Quantitative Measure of Acidity
Acidity is related to concentration of H+
(or H3O+)
pH = -log[H+]

7. pH Scale
In pure water, a few molecules ionize to form H3O+ and OH– H2O + H2O  OH– + H3O+
In acidic and basic solutions, these concentrations are not equal acidic: [H3O+] > [OH–] basic: [OH–] > [H3O+] neutral: [H3O+] = [OH–]

8. pH Scale Measure how much H3O+ is in a solution using pH
pH < 7.0 = acidic
pH > 7.0 = basic
pH = 7.0 = neutral
Measure of H3O+ and OH– concentration (moles per liter) in a solution
As acidity increases, pH ?

9. pH Scale
The pH scale is logarithmic: log(102) = log(101) = log(100) = –1 log(10–1) = – –2 log(10–2) = –2
pH = –log [H+]

10. pH Scale
pH = –log [H3O+]
What is the pH if [H+] = 1 x 10–5? Acidic or basic?
What is the pH if [H+] = 1 x 10–9? Acidic or basic?
What is the pH if [H+] = M?

11. pH of Strong Acid Solutions
For monoprotic strong acids: [H+] = [acid]
What is the pH of a 2.4 x 10-2 M solution of HNO3?

12. Finding [H3O+] from pH
[H+] = 10-pH
What is [H+] if pH = 8.9?