Chemistry 8 (Acids And Bases)

Front 1

Arrhenius Acid

Back 1

When dissolved in water, produces H3O+ or H+ (same thing)

Front 2

Arrhenius Base

Back 2

When dissolved in water, produces OH- ions

Front 3

Bronsted Lowry Acid

Back 3

H+ donor

Front 4

Bronsted Lowry Base

Back 4

H+ acceptor

Front 5

Lewis Acid

Back 5

Electron pair acceptor

Front 6

Lewis Base

Back 6

Electron pair donor

Examples: Ligands (coordinate covalent bonding), nucleophile

Front 7

Strong Acids (7)

Back 7

- H₂SO₄
- HCl, HBr, HI
- HNO3
- HClO₄
- HCO₃

Front 8

Strong Bases (3 Groups)

Back 8

- Group I and Group II Metal Hydroxides
- Group I Metal Oxides (Na₂O₂)
- Metal Amides (XNH₂)

Front 9

Periodic Trend for Acidity

Back 9

- Acidity increases diagonally right
- Radius increases diagonally left
- Increase radius increase force and increase charge

Front 10

Strong Acids do what in water?

Back 10

- Completely dissociate
- Produce stable conjugate base that does not react further in solution

Front 11

Acid Dissociation Constant

Back 11

- If > 1, strong. If < 1, weak.
- [Product]/[Reactant]

HA + Water --> A- + H+

[H+][A-]/[HA]

Front 12

Base Dissociation Constant

Back 12

- If > 1, strong. If < 1, weak.
- [Product]/[Reactant]

B- + Water --> OH + HB

[HB][OH-]/[B]

Front 13

Conjugates

Back 13

- The stronger the acid, the weaker its conjugate base
- The weaker the acid, the stronger its conjugate base**

** This is an exception... weak acids have a huge range. A weak acid like NH₄+ gives a weak base NH₃.

Front 14

pKa

Back 14

pKa = -logKa
pKb = -logKb

How do they relate?
↑ Ka ↓ pKa STRONG ACID
↓ Ka ↑ pKb WEAK ACID

Front 15

Auto-Ionization of Water

Back 15

Kw = [Products]/[Reactants]

- It's equal to 10e-14 at 25C. Varies with temperature.
- Ka * Kb = Kw

Front 16

pH

Back 16

- Measure of H+ in solution. Conversely, pOH is the measure of OH in solution.
- pH < 7 acidic, pH > 7 basic.

pH = -log[H+] and [H+] = 10^-pH

Front 17

How do we figure out pH/pKa of STRONG acids?

Back 17

- Strong acids completely dissociate in water
- Their pH is directly related to the molarity given (the power)
- Ex: 0.01M solution of HCl has a [H+] of 1e-2 and a pH of 2.

Front 18

How do we figure out the EXACT pH/pKa of STRONG acids and bases?

Back 18

- pH is always (n-1) the power of the molarity.

Eg: 3.43e-3M of HCl
- Figure out what it's between (2....3)
- Then subtract 3.43 from 10 = 6.57
- Move it over a decimal point and add it onto the lower number... voila!

Front 19

There are 14 equations for Acids/Bases. GO!

Back 19

1. pH = -log[H+]
2. [H+] = 10^-pH
3. pKa = -logKa
4. Ka = 10^-pKa
5. pKb = -logKb
6. pKa = -logKa
7. pKw = -logKw = 14
8. Kw = 10^-pKw = 10e-14
9. Kw = Ka*Kb = 10e-14
10.pKa + pKb = pKw = 14
11. pH + pOH = 14 = pKw
12. [H+] = [OH-] = 10e-7
13. pOH = -log[OH-]
14. [OH-] = 10^-pOH

Front 20

pH of Salts
Neutralization Equation

Back 20

- All salts occur in the same kind of reaction and the heat is the SAME for all neutralization reactions

Acid + Base --> Water + SALT + Energy(heat)

Front 21

Strong Acids and Strong Bases yield what kind of ions for salts?

Back 21

Neutral

SA --> Neutral ANION
SB --> Neutral CATION

Front 22

Weak Acids and Weak Bases yield what kind of ions for salts?

Back 22

WA --> Basic Anion
WB --> Acidic Anion

Front 23

How do you figure out the pH or Ka of a WEAK ACID?

Back 23

- Weak acids don't completely dissociate in water
- We use an ICE table here... Initial, Change, Equilibrium

Front 24

SHORTCUT to Weak Acids

Back 24

[H+] = √(Ka * [X])

Or use ICE table

Front 25

Henderson-Hasselbach Equation

Back 25

pH = pKa + log[A]/[HA]
pOH = pKb + log[HB]/[B]

Front 26

What is a Buffer?

Back 26

A buffer resists changes in PH

Front 27

What are the ingredients of a buffer?

Back 27

- WEAK ACID + salt of its conjugate base
- WEAK BASE + salt of its conjugate acid

Front 28

What makes a buffer good?

Back 28

- Buffer should have a pKa near the pH you want to maintain
- To make a really good buffer you use a lot of it

Front 29

Buffer Capacity

Back 29

- Use Henderson Hasslebach Equation
- Note that having a lot more of a buffer will resist a change a lot better than a smaller amount

Front 30

Indicators

Back 30

- Weak acid that undergoes a color change when its converted to its conjugate base
- Purpose is to show you the pH of a solution
- The color change occurs during deprotonation
- Indicator is usually pKa±1 near the pH we want to indicate (near equivalence point)

Front 31

Titrations: What do we use them for? (2)

Back 31

(1) Find concentration of a known acid
(2) Find identity of an unknown acid

Front 32

Titrant

Back 32

ALWAYS a strong acid and base and ALWAYS known

Front 33

Half-Equivalence Point

Back 33

- Point at which [HX] = [X-]
- Amount of acid equals the amount of its conjugate base

Front 34

Equivalence Point

Back 34

- Complete neutralization
- Point at which the mols of base = mols of acid originally present

Front 35

Buffering Region

Back 35

Section of the titration curve where the pH changes very gradually

Front 36

What can we find at the half-equivalence point?

Back 36

The pKa of our unknown. Since at this point the ratios are equal, we get pH = pKa according to Henderson-Hasslebach.

Front 37

Polyprotic Acid

Back 37

- Has more than one equivalence point
- Equivalence points is equal to the number of ionizable hydrogens the acid can donate
- Think Amino Acids