Chapter 9: Acids, Bases and Salts

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1

For a long time the concept of an acid was known. In
1661, Robert Boyle characterized an acid as

any substance
that when dissolved in water tastes sour, and was corrosive
to metal. A base (he used the term alkalies) was any
substance, that when dissolved in water tasted bitter, and felt
slippery to the touch. Note that the alkali metal group of the
periodic table contain metals that readily form bases.

2

In 1887, Swedish chemist, Svante Arrhenius

developed
the theory that an acid was any substance that when dissolved
in water produced protons, H+ and that a base was any
substance that when dissolved in water produced hydroxyl
ions, OH-. This theory is still useful and covers most of the common acids and bases, such as HCl, HNO3, H2SO4 and the
common bases NaOH, Ca(OH)2 and Mg(OH)2.

3

Bronsted Theory
Later, in 1923, Johannes Bronsted of Denmark and
Thomas Lowry in England

separately expanded the theory of
acids and bases to say that an acid was any hydrogencontaining
substance that could donate a proton when
dissolved in water. Thus an acid was a proton donor. The
base was any substance that could accept a proton, a proton
acceptor. This didn't so much add to the list of acids but it did
add to the list of bases and the concept of the acid base pair.

4

Bronsted
acid.

When a substance donates a proton it is called a Bronsted
acid. The species left behind is call the conjugate base. The
term conjugate acid-base pair is the Bronsted acid and its
conjugate base. The two differ by a single proton. The
conjugate acid has the proton, the conjugate base does not.

5

Identify the conjugate acid and bases.

HNO3 (aq) + H2O (l) ↔ H3O+(aq) + NO3 - (aq)
acid conj base ______ acid conj base
|_______________________________|

6

Identify the conjugate acid and bases.

NH3 (aq) + H2O (l) ↔ OH- (aq) + NH4 + (aq)
base conj acid ____ base conj acid
|_____________________________|

7

amphoteric

Notice that water can act as either the conjugate acid or the conjugate base depending on what ions are added to it.
This property of being able to both accept and donate a proton makes water amphoteric, which means having both forms.

8

Naming Acids

There are two sets of rules for naming acids. The first set is for acids consisting of a H atom bonded to a non-metal.
The second set is for acids containing a H atom bonded to a polyatomic ion, usually with oxygen.

9

HCl, HI, HBr, H2S

are all members of the first group.
Although each of these covalent molecules can be named as gases, in water they become acids and have different names

10

HCl -

hydrogen chloride becomes hydrochloric acid

11

HI

hydrogen iodide becomes hydroiodic acid

12

HBr

hydrogen bromide becomes hydrobromic acid

13

H2S

dihydrogen sulfide becomes hydrosulfuric acid.

14

The rules for naming these acids ( HCl, HI, HBr, H2S) are:

Drop word hydrogen and replace it with "hydro".
Add the root for the anion - chlorine goes to "chlor".
Add the suffix "ic" to the anion root
Add the word "acid" ==> HCl = hydrochloric acid.

15

The only exception in this series is

that H2O is water,
which can act as both an acid and a base - amphoteric. This is important !!

16

HNO3, H2SO4, H3PO4, H2CO3, HNO2, H2SO3 and H3PO3
are all members of the second group of acids

those that contain a polyatomic, oxygen-rich anion. In this case, the acid name comes from the name of the polyatomic anion.

17

The nitrate anion, NO3 -

forms nitric acid - HNO3.

18

The sulfate anion, SO4 2-

forms sulfuric acid - H2SO4.

19

The phosphate anion, PO4 3-

forms phosphoric acid - H3PO4.

20

The carbonate anion, CO3 2-

forms carbonic acid - H2CO3.

21

The nitrite anion, NO2 -

forms nitrous acid - HNO2.

22

The sulfite anion, SO3 2-

forms sulfurous acid - H2SO3.

23

The phosphite anion, PO3 3-

forms phosphorous acid - H3PO3

24

The rules for naming these acids( HNO3, H2SO4, H3PO4, H2CO3, HNO2, H2SO3 and H3PO3) Are :

Drop word hydrogen and do not replace it with anything.
Add the root for the anion - nitrate goes to "nitr".
Add suffix "ic" to the acid root if the anion ends in "ate".
Add suffix "ous" to the acid root if anion ends in "ite".
Add the word "acid" ==> HNO3. = nitric acid.

25

The Self-Ionization of Water

We have seen above that water can act as both a Bronsted acid and a Bronsted base depending on what other ions are in solution.
By itself, water will self-ionize to form
both hydronium and hydroxial ions:

26

neutral, H2O (l) + H2O (l) ↔ H3O+ (aq) + OH- (aq)

In pure water, at 25 0C, the concentration of both of these ions has been measured experimentally and found to be
present at 1x10-7 mol/L (M). From this comes the definition of neutral, which means the concentration of H3O+ and OH- are
equal. Starting with the equilibrium constant, Keq, for the formation of these water ions, we can derive a very useful constant called the water constant or ion product constant of
water, Kw, which has the value of Kw = 1x10-14 (mol/L)2

27

nuetral

Keq = [H3 O + ][ OH -] = [H3 O + ][ OH -]
[H2O][H2O] [H2O]2
Kw = Keq[H2O][H2O] = [H3O+ ][OH-]
Kw = [H3O+ ][OH-] = [1x10-7 mol/L][1x10-7 mol/L]
Kw = Keq[H2O][H2O] = 1x10-14 (mol/L)2
This constant is true not only for pure water, but for any water solution since most of the solution will be water molecules.

28

acidic

A solution is classified as being acidic if the concentration of H3O+ ions is greater than the concentration of OH- ions.

29

basic or alkaline

Conversely, the solution is consider to be basic or alkaline if the concentration of OH- ions is greater than the concentration of H3O+ ions. Also, one should remember that in either case,
the product of the H3O+ ions and the OH- ions will still be Kw.
This condition allows us to solve for the the hydroxyl concentration if given the hydronium concentration and vice versa.

30

Determine if the solution is acidic or alkaline
(basic) and what is the concentration of the other ion.

a) [H3O+ ] = 1x10-4 M => acidic, [OH-] = 1x10-14 M / 1x10-4 M = 1x10-10 M
b) [OH- ] = 1x10-9 M => acidic, [H3O+ ] = 1x10-14 M / 1x10-9 M = 1x10-5 M
c) [OH- ] = 1x10-6 M => basic, [H3O+ ] = 1x10-14 M / 1x10-6 M = 1x10-8 M

31

The pH Concept
In 1909, Soren Sorenson, a Danish chemist, introduced the pH scale as a simplified way of describing the
H3O+ ion concentration in solutions.

The mathematical definition of pH is the negative log of the molar hydrogen ion concentration.
pH = - log(H3O+)
Conversely the [H3O+ ] = H+ = 1x10-pH

32

PH Scale

The pH scale runs from 0-14 with numbers below 7 representing acidic solutions and numbers above 7 representing alkaline (basic) solutions. pH = 7 is neutral.

33

If the concentration of [H+] is given, then one can calculate the pH and the [OH-] concentration.

Determine the [H+] and the [OH-] gKeq =
[H3 O + ][ OH -] = [H3 O + ][ OH -]
[H2O][H2O] [H2O]2
Kw = Keq[H2O][H2O] = [H3O+ ][OH-]
Kw = [H3O+ ][OH-] = [1x10-7 mol/L][1x10-7 mol/L] Kw = Keq[H2O][H2O] = 1x10-14 (mol/L)2
given the pH.

34

Determine [H+] and [OH-] given the pH.

a) pH = 10.0, [H+] = 1x10-10 M => basic, [OH-] = 1x10-4 M.
b) pH = 4.0, [H+] = 1x10-4 M => acidic, [OH-] = 1x10-10 M.
c) pH = 5.0, [H+] = 1x10-5 M => acidic, [OH-] = 1x10-9 M.

35

If concentration is not a whole number, then one must use
logarithms to calculate pH.

[H+ ] = 3.6x10-4 M, then pH = -log(3.6x10-4 )= 3.44

36

Determine pH from [H+] concentration

a) [H+] = 4.2 x 10-5 M => acidic, pH = 4.38.
b) [H+] = 8.1 x 10-9 M => basic, pH = 8.09.

37

Determine [H+] concentration from pH.

a) pH = 2.75, acidic, [H+] = 10-2.75 = 1.78x10-3 .
b) pH = 8.33, basic, [H+] = 10-8.33 = 4.68x10-9 .

38

Properties of Acids

One of the properties shared by acids and bases is that they can be diluted from a stock solution to produce a required molar concentration.

39

One of the properties shared by acids and bases is that they can be diluted from a stock solution to produce a required molar concentration.

CcVc =CdVd ==> Vc =CdVd /Cc =
Vc =CdVd /Cc =(1 M HNO3)*(250 ml) /(6 M HNO3) = 41.7 ml.

40

Calculate how to make 500 ml of 3.0 M
aqueous ammonia from a concentrated stock
solution of 15 M NH3.
CcVc =CdVd ==> Vc =CdVd /Cc =

Vc=CdVd /Cc =(3.0 M HNO3)*(500 ml)/(15.0 M NH3)= 100 ml.

41

Another property common to all acids, besides tasting sour and producing H3O+ ions, is

that they react to form a double-replacement reaction with oxides, hydroxides,carbonates and bicarbonates.

42

Acids react to form a double-replacement reaction with oxides, hydroxides,carbonates and bicarbonates.

Because all acids react in this way, the important reaction
does not so much depend on the identity of the acid itself.
This gives rise to an abbreviated type of equation called a net
ionic equation, which only includes the ions of interest. All other ions are called spectator ions and are omitted.

43

To get to
the net ionic equation

one starts with the molecular equation,
then writes the total ionic equation to identify the spectator ions and after eliminating them, one is left with the net ionic equation.

44

Show molecular, total ionic and net ionic equations.

Mol eq. 2HCl (aq) + Ca(OH)2 (s) → CaCl2 (aq) + 2H2O (l) T.I. 2H+ (aq) + 2Cl- (aq) + Ca(OH)2 (s) → Ca2+ (aq) + 2Cl- (aq) + 2H2O(l) net ion eq. 2H+ (aq)+ Ca(OH)2 (s) → Ca2+ (aq)+ 2H2O(l) (Cl spectator)

45

Show molecular, total ionic and net ionic equations.

Mol eq. 2HCl (ag)+ Sr(HCO3)2 (s)→ SrCl2 (ag)+ 2CO2 (g) + 2H2O (l) T.I. 2H+ (aq)+ 2Cl-
(aq)+ Sr(HCO3)2(s) → Sr2+ (aq)+ 2Cl- (aq) + CO2(g) + 2H2O(l) net ion eq. 2H+ (aq)+ Sr(HCO3)2 (s)→Ca2+
(aq)+CO2 (g)+2H2O(l) (Cl spect)

46

Another property of acids is their ability to react with
metal,

to varying degrees, to form a metal salt (is there another kind?) and hydrogen gas in a redox reaction. The metal is being oxidized from the pure state to form the salt. The hydrogen of the acid is reduced to form pure hydrogen gas.
Different metals react with acids at different rates forming an activity series from potassium as most reactive to gold and other noble metals as non-reactive.

47

Properties of Bases In 1661, Boyle said that bases not only feel slippery
but they also react to reduce the strength of an acid.

This has become known as a neutralization reaction, which if allowed to go to completion, the acid and base will consume each other
to produce only a salt and water. The classic example is: hydrochloric acid + sodium hydroxide yields sodium chloride and water Mol eq. HCl (aq) + NaOH (s) → NaCl (aq) + H2O (l)
T.I. H+ (aq) + Cl- (aq) + Na+ (aq) + OH- (aq) → Na+ (aq) + OH- (aq) + H2O (l) Net ion eq.H+ aq) + OH- (aq) → H2O (l) (Na+ (aq) and Cl- (aq) spectator)

48

Arrhenius bases are a metal salt containing the hydroxide
anion. These include:

Barium hydroxide (Ba[OH]2)
Calcium hydroxide (Ca[OH]2)
Lithium hydroxide (LiOH)
Potassium hydroxide (KOH)
Sodium hydroxide (NaOH)
Strontium hydroxide (Sr[OH]2)

49

Bronsted Lowry bases include all the Arrhenius bases and
add those which are related to ammonia. These include:

Ammonia (NH3)
Primary amines like : Methylamine (CH3NH2)
Secondary amines like : dimethylamine (CH3NHCH3)
many other organ compounds.

50

Salts

One way of thinking of a salt is as a crystalline, ionic compound that contains the positively charged cation of a base and the negatively charge anion of an acid.

51

There are a number of ways to form a salt. We will look
at three ways:

1) acid + metal → salt + hydrogen gas
2) acid + metal oxide → salt + water
3) acid + metal hydroxide → salt + water

52

Show three ways to form Mg(NO3)3.

1) 2HNO3 (aq) + Mg (s) → Mg(NO3)2 (aq) + H2 (g)
2) 2HNO3 (aq) + MgO (s) → Mg(NO3)2 (aq) + H2O (l)
3) 2HNO3 (aq) + Mg(OH)2 (s) → Mg(NO3)2 (aq) + 2H2O (l)

53

Show three ways to form AlCl3.

1) 6HCl (aq) + 2Al(s) → 2AlCl3 (aq) + 3H2 (g)
2) 6HCl (aq) + Al2O3 (s) → 2AlCl3 (aq) + 3H2O (l)
3) 3HCl (aq) + Al(OH)3 (s) → AlCl3 (aq) + 3H2O (l)

54

Strengths of Acids and Bases strong electrolytes

When salts dissolve in water, they usually do so
completely and are called strong electrolytes.

55

weak
electrolytes.

Some salts
only dissolve poorly in water and are called weak
electrolytes.

56

Not all acids or bases dissolve fully in water.

Those that do are called strong acids or strong bases. Those
that do not fully dissolve are called weak (moderately weak)
acids or weak bases. The dissolution of salt or an acid is a
reversible reaction although in the case of a strong electrolyte
or a strong acid (or base) the forward reaction is nearly 100%.
The case of a weak electrolyte or a weak acid (or base) the
forward reaction can be nearly absent.

57

The strength of an acid (or base) is determined

by the size
of its own acid dissociation constant, Ka, which equals the
ratio of the hydronium ion concentration times the conjugate
base concentration over the undissociated acid concentration.