Chapter 14

Acids and Bases

Acid/Base Theories

• Arrhenius Theory– Acids produce H+ ions in solution– Bases produce OH- ions in solution– Downside

• Must be in solution and must have those ions

• Bronsted-Lowry Theory– Acids are H+ donors (Proton donors)– Bases are H+ acceptors (Proton acceptors)

Vocabulary

• H+ is the hydrogen ion– Just a proton

• H3O+ is a hydronium ion

– It is the way H+ exists in water

• Water accepts a hydrogen ion and becomes H3O+

• Either way is fine the first is just easier

Conjugates Acid/Base Pairs

• Conjugate Acid is formed when a base gains a proton

• Conjugate Base is what remains after an acid donates a proton

• Ex – HNO3 + H2O H3O+ + NO3-

General Form for Acids and Bases

HA + B A- + BH+

• HA is an acid

• B is a base

• A- is the conjugate base– Just the negative ion of the acid

• BH+ is the conjugate acid– Just the base plus a hydrogen

Strong and Weak Acids

• Strong acids completely ionize in solution– Nitric, Perchloric, Sulfuric, Hydrochloric,

Hydrobromic, Hydroiodic– Weak Conjugate bases

• Weak Acids only partially ionize in solution– Every other acid– Equilibrium is established in the ionization– Weak acids have Ka values – Strong Conjugate bases

Acid Dissociation Constant, Ka

HA(aq) + H2O A- + H3O+

• For weak acids equilibrium is established

• Equilibruim constant is Ka

• Ka=[A-][H3O+]/[HA]

• Values tend to be small– Because CB is fairly strong

• Strong acids do not have Ka values.

Acid Terms

• Monoprotic – One acidic hydrogen

• Polyprotic – Many acidic hydrogens

• Diprotic – Two acidic hydrogens

• Triprotic – Three acidic hydrogens

• Oxyacid – Acid that has an acidic hydrogen attached to an oxygen

• Organic Acid – Acid that has the acidic hydrogen attached to the carboxyl group

Water As An Acid and Base

• Water is amphoteric – Both an acid and base

H2O + H2O H3O+ + OH-

• Equilibrium system that always has the same value

• Called Kw

Autoionization of Water, Kw

H2O + H2O H3O+ + OH-

• Kw = [H3O+] [OH-]

• In pure water the products have the same concentration, 1.00x10-7M

• Value of Kw = 1.00x10-14 at 25ºC

• Concentrations can change is acid or base is added

Acid, Base, or Neutral

• If the concentration of H+ = OH-

– Neutral

• If the concentration of H+ > OH-

– Acidic

• If the concentration of H+ < OH-

– Basic

What is the hydrogen ion concentration when the hydroxide ion concentration is 1.00x10-5M

What is the hydroxide ion concentration when the when the concentration of nitric acid is 0.0010M?

Homework

• P. 704 30, 32, 33, 35, 39ab,40ab

Logarithms

• The logarithm of a number to a given base (commonly 10) is the power or exponent to which the base must be raised in order to produce the number.

• SAY WHAT!

Examples

• If your question says

log 100 = x

It is saying to what power must 10 be raised to equal 100

10x = 100

102 = 100

So x = 2

Examples

log 1 = x x = 0

log 10 = xx = 1

log 1000 = x

x = 3

log 1x106 = x

x = 6

Examples

• Logarithms can also be used for numbers smaller than 1

log 0.1 = x

x = -1

log 0.01 = x

x = -2

log 1x10-5

x = -5

Examples

• If they are not easy to calculate you can do it on your calculator

log 15 = x

You can approximate it between . . .

–Type log 15 on your calculator

x = 1.18

Examples

• If your question says

log x = 7

It is saying 10 to the 7th power is what number

107 = 1x107

log x = 3

x = 1000

log x = -3

x = 0.001

Tougher Examples

log 234 = x

x = 2.37

log x = -3.3

x = 5.0x10-4

-log 9.1x10-5 = x

x = 4.0

-log x = 12.1

x = 7.9x10-13

pH

• Negative logarithm of the concentration of hydrogen ions in solution

pH = -log [H+]

• pH means power of Hydrogen

• Measures how acidic or basic a solution is

• pH scale typically goes from 0 to 14

• pH < 7 Acidic

• pH > 7 Basic

• Highly acidic = low pH

• Highly basic = high pH

Significant Figures and pH

• Digits after the decimal are the only ones that are significant in pH values

pH = 4.44

2 Significant Figures

pH = 10.874

3 Significant Figures

• If your [H+] is 0.088 M your pH is 1.06

Determining [H+] in solution

• The concentration of a strong acid is equal to the H+ concentration.

• 0.010 M HCl has an [H+] of 0.010 M

• To obtain the [H+] for weak acids you must use equilibrium

– Need Ka data

• Discuss hydroxide later

Other Info

Turn Kw into a log equation

pH + pOH = 14.00

pH’s You MUST Know

• When the [H+] is _______ the pH is _____0.10 M

1.000.0010M

3.001.0x10-6M

6.001.0x10-10M

10.00

• Find the pOH, [H+], [OH-] of lemon juice that has a pH of 2.48

• Determine the pH of 0.150M HCl

• Determine the pH of 2.3x10-3 M Hydrocyanic Acid HCN. Ka = 6.2x10-10

Homework

• Page 705 #’s 45,47,50,52,53,58

Mixtures of Weak Acids

• When there are mixtures of weak acids in solution determining the pH could be a difficult problem.

• However

– The acid with the largest Ka will control the pH of the solution

• Solutions of 0.10M HF (Ka = 7.2x10-4) and 0.10M HCN (Ka = 6.2x10-10) are mixed A) Which acid will control the pH of the solution? Why? B) What is the pH of the resulting solution.

Percent Dissociation

• Ratio of the concentration of the dissociated ions to the initial concentration

• Found just like the doing the 5% check

• Can be used to find pH and Ka

100*%centrationInitialCon

rationIonConcentHonDissociati

• A 0.25M solution of HClO is 20.% dissociated. A) What is the pH? B) What is the Ka value of the acid?

Strong Bases

• Strong bases are any compound containing the hydroxide ion

– Group 1 hydroxides are very soluble

– Group 2 less soluble but still strong

• Group 2 have two hydroxides per mole

–Be Careful

• Determine the pH of 0.022M Sr(OH)2

Weak Bases

• Organic Bases are weak bases. (Ammonia)

– Contain Nitrogen (Amines)

• CH3NH2 – Methylamine

– Must have a lone pair of electrons

– The hydroxide will come from water

(CH3)3N + H2O

Cont.

• Weak bases have Kb values

B + H2O BH+ + OH-

• The conjugate acid of a weak base is stronger than water.

][

]][[

B

OHBHKb

• Determine the pH of 15.0 M ammonia. The Kb is 1.8x10-5

Homework

• Page 705 #’s 62,63a,72,73,77ab,87

Polyprotic Acids

• Acids with more than one acidic hydrogen• Dissociate in a “stepwise” process

H3PO4 H+ + H2PO4- Ka1 =7.5x10-3

H2PO4- H+ + HPO4

-2 Ka2 =6.2x10-8

HPO4-2 H+ + PO4

-3 Ka3 =4.8x10-13

• Ka1 > Ka2 > Ka3

• Successive dissociation do not effect pH (except sulfuric acid)

Determine the pH of 5.0M H3PO4 (Ka=7.5x10-3) and the [H2PO4

-], [HPO4-2], [HPO4

-3]

Sulfuric Acid

• Sulfuric acids has two dissociations

– The first is strong

– The second is weak

• The second dissociation is quite strong, but it is not complete

H2SO4 H+ + HSO4- (Strong)

HSO4- H+ + SO4

-2 Ka = 1.2x10-2

Continued

• When you write sulfuric acid in net ionic equations only use the first equation

H2SO4 H+ + HSO4-

• The second dissociation only needs to be considered in dilute solutions.

• A 1.0M solution of sulfuric acid will have a lower pH than a 1.0M solution of HCl

Acid / Base Properties of Salts

• Some salts have acid base properties

– Make a solution have a pH below or above 7

• Some have no acid base properties

– Make a solution with a pH of 7

Neutral Salts

• Contain the following

• A metallic ions– (Except Aluminum)

• A conjugate base of a strong acid– Such AS . . . .

• Conjugate base of strong acids are weaker than water.

Basic Salts

• Contain the following

• A metallic ion

• A conjugate base of a weak acid– Such AS . . .

• Conjugate base of weak acids are stronger than water.

Acidic Salts

• Contain the following

• A conjugate base of a strong acid

• A conjugate acid of a weak base– Such AS . . .

• Conjugate acid of weak bases are stronger than water.

Acidic Aluminum Ions

• Highly charged cations polarizes the O-H bond in water

• Hydrogens in water become acidic

• Ion becomes hydrated

• Al(H2O)6+3

OMGosh I Have Both Types of Ions

• If a salt has acidic and basic ions – Such AS . . .

• If Ka > Kb the solution is acidic

• If Ka < Kb the solution is basic

• If Ka = Kb the solution is neutral

Converting Ka to Kb

• If you know an acids Ka value you can find its Kb value as a salt.

• If you know a bases Kb value you can find its Ka value as a salt

Kw = Ka * Kb

Calculate the pH of 0.33 M NaHCO3

Homework

• P 707 #’s 94,95b, 98,99,102,105a, 107,112

Effect of Structure On Acids

• Different bonding patterns lead to differences in the strength of acids

• Two different types of acids

– Hydrohalic

– Oxyacids

Hydrohalic Acids

• Acids that contain a hydrogen and a halogen

• The weaker the bond the stronger the acid

Bond Enthalpy

H—F 567

H—Cl 433

H—Br 366

H—I 299

Oxyacids

• Tend to be of the form H – O – X

• The greater the electronegativity of X the stronger the acid

• Pulls electron density away and weakens the H – O bond

• Which is a stronger acid?

• H – O – Cl or H – O – Br – H – O – Cl because Cl has a greater E.N.

Oxyacids

• Bases tend to be of the form X – O – H too

• NaOH

• Why

• Metal has a low electronegativity and the O – H bond is not weakened

Oxyacids

• Acids with the same X element can have different number of oxygens

• As the number of oxygens increases so does the strength of the acid

• The extra oxygens weaken the H – O bond

Oxides

• Nonmetalic oxide in water gives an acid

CO2 + H2O H2CO3

• Acidic anhydrides (Acid w/o water)

• Nonmetal must keep its oxidation #

• Metallic oxides in water gives a base

Na2O + H2O 2NaOH

• Basic anhydrides

Lewis Acids and Bases

• Lewis acids are electron pair acceptors

• H+ for example has no electrons

• Lewis bases are electron pair donors

• NH3 has an electron pair to share

Homework

• P 708 #’s 113,114,115ab,116ab,119