1451ch15

Chapter 15: Acids and Bases

Key Topics:

Chapter 15 builds on your knowledge of equilibrium and ICE charts and applies what you learned to Acids and Bases. That being said, if you are still struggling with ICE charts, you need to get that chapter hammered out before getting too deep into this chapter.
This Chapter begins by describing the characteristics of an acid and a base and looking at them from both the Arrhenius and the Bronsted-Lowry definitions.
- Acids
  - are sour in taste, and are often used to clean metals.
  - The Arrhenius definition of an acid states that an acid is a molecule that dissociates in water to produce H+.
  - The Bronsted-Lowry definition of an acid states that an acid is a molecule that serves as a proton (H+) donor. This definition is more broad as it is not specific to an aqueous solution. Note that everything that is defined as an acid using the Arrhenius definition is still an acid under this new definition as well.
- Bases
  - are bitter in taste, and feel slippery.
  - The Arrhenius definition of a base is a compound that dissociates in water to produce OH-.
  - The Bronsted-Lowry definition of a base is a substance that is a proton acceptor. This widens the definition by including molecules that have lone pairs that can be used to bond with an H+. The most common base that is a B-L base but not an Arrhenius base is ammonia (NH3). Again, just like the acids, everything that is considered a base by the Arrhenius definition is still a base under the B-L definition.
pH
- pH is a measure on a log scale of the H3O+ concentration. It is calculated as -log(H3O+) and the lower the pH, the more H3O+ is in the solution.
- You can also calculate the pOH,the pKa, or the pAnything by taking the -log of the value.
- Using the water self reaction, we can see that the pH + pOH = 14 for any solution (even after acid or base is added). We also see that 10^-14 = [H3O+] [OH-]. This means that there is always some OH- in a solution -even if the solution is very acidic! It also means that the pH of pure water is 7.
Acid Strength
- Strong acids dissociate completely and thus can use a one-sided arrow when looking at their reaction with water: HA(aq) + H2O(l) --> H3O+(aq) + A-(aq).
  - When considering a solution that contains a single strong acid, the H3O+ and the A- concentrations are equal to the initial concentration of the acid (HA).
- Weak acids do not dissociate completely, so you must use an ICE Chart and equilibrium constant to calculate the concentrations of H3O+ and A-. These equilibrium constants are given a special name (Ka), but still work the exact same as any other K.
- Larger values of Ka mean stronger acids.
- Smaller values of pKa mean stronger acids
- Percent ionization is a measure of much of the acid has dissociated in the solution. As you increase the concentration of a weak acid, you increase the overall [H3O+], but decrease the %ionization.
Calcualtions: Acids
- You should be able to calculate the pH, and all ion concentrations for the following solutions:
  - Solutions of a single Strong acid
  - Solutions of a single Weak Acid
  - Solutions of a strong acid mixed with a weak acid
  - solutions of 2 weak acids
  - Polyprotic acids (treat these like a SA/WA or WA/WA mixture)
Calculations: Bases

You should be able to write a reaction between a base and water
Use a Kb to describe a base's strength.
You should be able to calcualte the pH, and all ion concentrations for the following solutions:

Solution containing a single strong base
Solution containing a single weak base
mixtures of Weak bases
A strong base mixed with a weak base.

Acid/Base properties of salts

Be able to predict whether a solution containing an ionic salt will be acidic/basic/or neutral.

CHAPTER 16 (Part 1)

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