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Chemistry – April 21, 2017

P3 Challenge –

Consider: 2 SO2 (g) + 1 O2 (g)  2 SO3 (g)

Determine which way the equilibrium will shift when the following disturbances
are applied.

A) SO3 is removed.

B) Temperature is increased

C) Volume is increased.
Get out LeChatelier’s Principle
worksheet for HMW check
.
(Exothermic) at equilibrium.
Chemistry – April 21, 2017

Today’s Objective –


Acid/Base equilibrium and pH
Assignment:

Acid/Base pH Worksheet

Agenda

Acid/base properties

Acid/base definitions

Acid/base concentrations

Dissociation and Ionization

Self-ionization of water

pH, pOH
Already about acids and bases (U5)

Acids react with bases to form water and a salt.

Type of Double Replacement reaction called a neutralization.

Forming water is a driving force

Acidic compounds start with H and have 3 naming rules.

There are 6 strong acids: HCl, HBr, HI, H2SO4, HNO3 and HClO4

All other acids are weak.

Strong bases are Group 1 and 2 hydroxides.
Acid and base properties

Acids properties:
 Sour
taste
 Dissolves
 Turns
metals

Base properties
 Bitter
taste
 Slippery
feel
litmus paper red
 Turns
litmus paper blue
phenolphthalein
colorless
 Turns
phenolphthalein
 Turns
pink
Arrhenius Definitions

Proposed by Svante Arrhenius, a Swedish scientist, in 1850

Acid = produces H+ ion in solution

Base = produces OH- ion in solution

Can label any substance based on its chemical formula.

Ex: NaOH

Ex: H2SO4

Limited usefulness because it misses some substances that we can
label as acids and bases based on their properties: NH3 is a base,
NaHCO3 is a base
Brønsted-Lowry Definition

Proposed independently by Johannes Nicolaus Brønsted and
Thomas Martin Lowry in 1923

Acid = proton (H+) donor

Base = proton (H+) acceptor

Some molecules can both donate and accept protons (Ex: Water)
and are called amphoteric.

Ex: NH3(aq) + H2O (l)  NH4+ (aq) + OH- (aq)

Ex: HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq)
Types of acids

Monoprotic acids donate 1 proton


Diprotic acids can donate 2 protons


Ex: HCl
Ex: H2SO4
Triprotic acids can donate 3 protons

Ex: H3PO4
Conjugate acids and Conjugate bases

The reactions of acids and bases are often reversible creating an equilibrium

The “acid product” that forms from the addition of a proton to a base is a
conjugate acid.

The “base product” that forms from the removal of a proton from an acid is a
conjugate base.

Any reactant and its conjugate partner form a conjugate acid-base pair.

Ex: NH3(aq) + H2O (l)  NH4+ (aq) + OH- (aq)
Dissociation vs Ionization


When an ionic compound dissolves in water =
dissociation.

The ions already exists, they just move apart.

Ex: NaOH a strong base added to water is a dissociation.
When a molecular compound splits apart and forms ions =
ionization

Ions must be created.

Ex: HC2H3O2 a weak acid, added to water is a partial ionization.
Strong Acid Concentrations

6 Strong acids completely ionize in water and are strong
electrolytes.

HCl, HBr, HI,

HNO3, HClO4, H2SO4

All other acids are weak, only partially ionize. Form an equilibrium.

HA (aq)  H+ (aq) + A- (aq)

For strong acids, determine ion concentrations to estimate [H+]

1.5 M HCl

3.5 M H2SO4
Strong Base Concentrations

Strong bases are the hydroxides of alkali and alkaline earth metals.

Strong bases are strong electrolytes and dissolve completely in
water.

All other bases are weak and only partially dissolve.

Determine the [OH-] concentration for

0.45 M Ba(OH)2
Water

Because water is amphoteric: can act as an acid or a base.

Water can react with itself, acting as both the acid and base:
H2O(l) + H2O(l)  H3O+ (aq) + OH- (aq)

This reaction is the self-ionization of water, and happens in very
small amounts.

In pure water at 25◦C, [H+] = [OH-] = 1 x 10-7 M = 0.0000001 M

Product of these two concentrations is called the ion product of
water
Kw = 1 x 10-14 = [H+] [OH-]

The ion product of water is true for all acidic or basic solutions.
H3O+ = H+
Relative concentrations
[H+]
1.E-01
1.E-02
1.E-03
1.E-04
1.E-05
1.E-06
1.E-07
1.E-08
1.E-09
1.E-10
1.E-11
1.E-12
1.E-13
[OH-]
1.E-13
1.E-12
1.E-11
1.E-10
1.E-09
1.E-08
1.E-07
1.E-06
1.E-05
1.E-04
1.E-03
1.E-02
1.E-01
Kw
1.E-14
1.E-14
1.E-14
1.E-14
1.E-14
1.E-14
1.E-14
1.E-14
1.E-14
1.E-14
1.E-14
1.E-14
1.E-14
pH
[H+]
1.E-01
1.E-02
1.E-03
1.E-04
1.E-05
1.E-06
1.E-07
1.E-08
1.E-09
1.E-10
1.E-11
1.E-12
1.E-13
Opp of Exp
1
2
3
4
5
6
7
8
9
10
11
12
13
pH
1
2
3
4
5
6
7
8
9
10
11
12
13

Easier to refer to [H+] based on the
exponent.

More specifically the opposite of
the exponent because they are
negative.

We call this value pH.

Small pH values, have highest [H+]

Mathematically,

pH = -log [H+]
Acidic, Basic, or Neutral pH

pH is by far the most common way to describe the
acidity/basicity of a solution

Always remember that pH is a log function: each
increase by 1 is a factor of 10.

Ex: pH = 7
pH = 3.2
pH = 8.2
pOH

It is less common than pH, but you can similarly define pOH
pOH = -log [OH-]

Low pOH values are basic, high pOH values are acidic, 7 is neutral.

pH +pOH = 14

Ex: If pOH = 4, pH =

Ex: If pOH = 10, pH =

Often easiest, to convert pOH to pH when identifying acidity.
pH Calculations

Four related quantities: pH, pOH, [H+] and [OH-]

pH = -log [H+] , pOH = -log [OH-], Kw = [H+] [OH-] = 1 x 10-14 , pOH + pH = 14

Note: the opposite of pH = -log [H+] is [H+] = 10 –pH

Calculate the other three quantities for the given information. Then indicate if it
is an acidic or basic solution.

[H+] = 4.0 x 10-3 M

[OH-] = 1.5 x 10-5 M

pH = 2.1

pOH = 3.5
Exit Slip - Homework

Exit Slip: An acidic solution has [H+] = 4.5 x 10-4 M. What are
a) the pH
b) the pOH and c) [OH-]

What’s Due? (Pending assignments to complete.)


Acid/Base pH worksheet
What’s Next? (How to prepare for the next day)

Read p496 - 520