Chapter 18: Aqueous Ionic Equilibrium (Exam 3) Flashcards Preview

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Flashcards in Chapter 18: Aqueous Ionic Equilibrium (Exam 3) Deck (30)
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1

Buffer

- Resists pH change by neutralizing added acid or base

1) Significant amounts of both weak acid and conjugate base

2) Significant amounts of both weak base and conjugate acid

2

Common Ion Effect

Calculate pH?

- When solution contains two substances that share common ions, solution ionizes less than it normally would

- Write balanced equation, subtract x from reactants and add x to products; substitute into equilibrium concentrations, use x is small approximatio, and determine H3O+ from calculated value of x 

- Substitue into pH equation pH = log[H3O+]

3

Henderson-Hasselbach Equation

- pH = pKa + log [base]/[acid]

- Remember, pKa = -logKa

- Use when x is small

4

Generic Approach to Calculating pH in Buffer Solution

1) Stoichiometry calculation in which we calculate the addition changes relative amounts of acid and conjugate base.

2) Equilibrium calculation in which we calculate pH based on new amounts of acid and conjugate base.

--> Stoichiometry: As added acid is neutralized, it converts the stoichiometric amount of base into conjugate acid through neutralization reaction

5

1) Effect of adding small amount of strong acid to buffer solution?
2) Effect of adding small amount of strong base to buffer?

1) Converts stoichiometric amount of base to conjugate acid and LOWERS pH of the buffer.

2) Converts stoichiometric amount of acid to conjugate base and RAISES pH of the buffer.

6

What does an effective buffer accomplish?

- Neutralizes small to moderate amounts of added acid or base

7

What is a buffer's effectiveness influenced by?

- Relative amounts of acid and conjugate base

- Absolute concentrations of acid and conjugate base

- Most effective when concentration of acid/base are equal and high; range must be 0.10-10 

8

How to prepare a buffer?

- Calculate ratio of conjugate base to acid required to attain desired pH

- Plug in pH, pKa, and use [base]/[acid] = 10n

- where "n" is equal to pH - pKa

- (Use H-H equation)

9

Buffer Capacity

- Amount of acid or base that can be added to a buffer without changin the pH significantly 

 

10

Acid-Base Titration

- Basic (or acidic) solution of unknown concentration reacts with acidic (or basic) solution of known concentration

11

Indicator

- Substance whose color depends on pH; monitors titration

12

Equivalence Point

- Point in titration in which number of moles of base is stoichiometrically equal to number of moles of acid

13

Titration: Strong Acid-Strong Base

- Before equivalence point, H3O+ is in excess

- Calculate [H3O+] by substracting number of moles added of OH- from initial moles of H3Oand dividing by total volume

- Equivalence point is pH = 7

- Beyond equivalence point, OH- is in excess

14

Titration: Weak Acid-Strong Base

- Calculate initial

- Calculate half equivalence point

- Calculate beyond equivalence point

- pH at equivalence point

- Initial pH is that of weak acid solution to be titrated 

- Between iitial pH and equivalence point, solution becomes a buffer (calculate with reaction stoichiometry to calculate amounts of each buffer component)
- Halfway to equivalence point: Buffer components are exactly equal (pH = pKa) --> (calculate: equilibrium problem for ionization of water by ion acting as weak base; # moles ion / volume at equivalence point)
- Beyond equivalence point, OH- is in excess (calculate: OH- by subracting initial moles weak acid from number moles added OH- and divide by total volume)

- Equivalence point: pH = 8

15

Titration: Weak Base-Strong Acid

- Calculate pH?

- Equivalence point?

- Use Henderson-Hasselbach equation

- pKa = conjugate acid being titrated

- Equivalence point: pH = 5.26

16

Titration: Polyprotic Acid

- pH curve has two equivalence points

- Volume required to reach first equivalence point is identical to volume required to reach second

17

Endpoint

- Occurs at equivalence point; point in which indicator changes color

- Indicators are typically weak organic acids

18

Solubility-Product Constant (Ksp)

- Equilibrium constant for a chemical equation representing the dissolution of an ionic compound

- Meausre of solubility of a compound

- Written as concentration of products [C]c[D]d

A + B --> cC + dD

19

Molar Solubility

- Solubility in units of moles/liter (calculated from Ksp)

- Calculate with ICE box

20

Effect of pH on solubility

- Solubility of an ionic compound with strongly basic or weakly acidic anion increases with increasing acidity

21

Q < Ksp

- Solution is unsaturated; more solid ionic compound can dissolve in solution

22

Q = Ksp

- Solution is saturated; holds equilibrium amount of dissolved ions

- Added solid does not dissolve ins olution

23

Q > Ksp

- Solution is supersatruated 

- Under most circumstances, excess solid precipitates out of supersaturated solution

24

Predicting Precipitation Reactions

- Use reaction quotient to compare to Ksp

25

Selective Precipitation

- Process involving the addition of a reagent that forms a precipitate with one of the dissolved cations (but not the others)

26

Complex Ion

- Contains central metal ion bound to one or more ligands

27

Ligands

- Neutral molecule or ion that acts as a Lewis base with central metal ion

28

Formation Constant

- Kf

- Equilibrium constant associated with reaction for the formation of a complex ion

- Written the same as a reaction quotient

29

Effect of Complex Ion Equilibrium on Solubility

- Solubility of an ionic compound containing a metal ion increases in the presence of Lewis bases that form complex ions with the cation

30

Effect of solution pH on solubility of amphoteric metal hydroxides?

- Metal hydroxides become more soluble in acidic solutions; can acts as base and react with H3O+

- Some metals can also act as an acid (amphoteric)