•Naming Acids…Slide 3
• Acids,
Bases and Neutralization Reactions
…Slide 8
•Calculation of pH…Slide 14
•Strength of Acids and Bases …Slide 23
• C.7.A name, acids using International Union of Pure
and Applied Chemistry (IUPAC) nomenclature rules
Naming Acids without Oxygen
 Acids without Oxygen are named with the prefix
“Hydro” and end in “-ic”
 Examples:
 HCl
Hydrochloric Acid
 HF
Hydrofluoric Acid
 HBr
Hydrobromic Acid
Naming Acids with Oxygen
 For acids with oxygen in several forms, prefixes are
used with the regular “-ic” and “-ous” endings.
 The “-ic” or regular ending for an acid comes from the
polyatomic ion with the “-ate” ending. This gives the
regular count for the oxygen for this type of acid.
 Example:
 H2SO4
 SO4 is Sulfate so this acid is called Sulfuric Acid
 Once you know the “-ic” ending, count the number of
oxygens in the other forms to find the name of the acid.
(REMEMBER: The regular “-ic” form comes from the
polyatomic ion that ends with “-ate”)
 Two less oxygen Hypo ________ “-ous” Acid
 One less oxygen
________ “-ous” Acid
 Regular “-ic” form
________ “-ic” Acid
 One more oxygen Per ________ “-ic” Acid
 The other names for the acids will come from the
count based from the “regular acid name”
 H2SO4 “-ate” ending so it is Sulfuric Acid
 H2SO3 “-ite” ending so it is Sulfurous Acid
 H2SO2 two less oxygen will have a prefix and
“-ous”ending. Hyposulfurous Acid.
 H2SO5 one more oxygen will have a prefix “Per”
and the regular “-ic” ending. Persulfuric Acid
C.10.G define acids and bases and distinguish between
Arrhenius and Bronsted‐Lowry definitions and predict
products in acid‐base reactions that form water
C.10.H understand and differentiate among
acid‐base reactions, precipitation reactions, and
oxidation reduction reactions
Acids: General Information
 Acid: a substance which when added to water
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produces hydrogen ions [H+].
Properties:
Acids: react with zinc, magnesium, or aluminum and
form hydrogen gas (H2).
react with compounds containing CO32- and form
carbon dioxide and water.
turn litmus red.
taste sour (lemons contain citric acid, for example).
DO NOT TASTE ACIDS IN THE LABORATORY!!
Bases
 Base: a substance which when added to water
produces hydroxide ions [OH-].
 Bases: feel soapy or slippery
 turn litmus blue
 they react with most cations to precipitate
hydroxides
 taste bitter (ever get soap in your mouth?) DO
NOT TASTE BASES IN THE LABORATORY!!
Arrhenius Model
 Basis for the model--action in water
 acid definition: produces H+ in water solution
 base definition: produces OH- in water solution
Bronsted-Lowry Model
 Basis for the model-- proton transfer
 acid definition: donates a proton
 base definition: accepts a proton
 conjugate acid definition: the acid becomes the
conjugate base after it donates the proton because it
can now accept it back.
 conjugate base definition: the base becomes the
conjugate acid after it accepts the proton because it
can now donate it back.
Acid-Base Reactions
 Strong acid + strong base: HCl + NaOH → NaCl +
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H 2O
net ionic reaction: H+ + OH- → H2O
Strong acid + weak base:
example: write the net ionic equation for the reaction
between hydrochloric acid, HCl, and aqueous
ammonia, NH3.
· Strong base + weak acid:
example: write the net ionic equation for the reaction
between citric acid (H3C6H507) and sodium hydroxide.
C.10.I define pH and use the hydrogen or hydroxide ion
concentrations to calculate the pH of a solution
Water Dissociation:
 Water dissociation: H2O(l) → H+(aq) + OH-(aq)
 equilibrium constant, KW = [H+][OH-] / [H2O]
Note: water is not involved in the equilibrium
expression because it is a pure liquid, also, the amount
of water not dissociated is so large compared to that
dissociated that we consider it a constant
 Value for Kw = [H+][OH-] = 1.0 x 10-14
 [H+] for pure water = 1 x 10-7
[OH-] for pure water = 1 x 10-7
 Definitions of acidic, basic, and neutral solutions
based on [H+]
 acidic: if [H+] is greater than 1 x 10-7 M
basic: if [H+] is less than 1 x 10-7 M
neutral: if [H+] if equal to 1 x 10-7 M
Example 1:
 What is the [H+] of a sample of lake water with [OH-]
of 4.0 x 10-9 M? Is the lake acidic, basic, or neutral?
 Solution: [H+] = 1 x 10-14 / 4 x 10-9 = 2.5 x 10-6 M
 Therefore the lake is slightly acidic
 Remember: the smaller the negative exponent, the
larger the number is.
 Therefore:
 acid solutions have exponents of [H+] from 0 to -6.
basic solutions have exponents of [H+] from -8 to -14.
Example 2:
 What is the [H+] of human saliva if its [OH-] is 4 x 10-8
M? Is human saliva acidic, basic, or neutral?
 Solution: [H+] = 1.0 x 10-14 / 4 x 10-8 = 2.5 x 10-7 M
 The saliva is pretty neutral.
pH
 relationship between [H+] and pH
 pH = -log10[H+]
 Definition of acidic, basic, and neutral solutions based
on pH
 acidic: if pH is less than 7
basic: if pH is greater than 7
neutral: if pH is equal to 7
 The [H+] can be calculated from the pH by taking the
antilog of the negative pH or 10-pH.
 Example 3: calculate the [OH-] of a solution of
baking soda with a pH of 8.5.
 Solution: First calculate the [H+]
 at pH 8.5, the antilog of -8.5 (or 10-8.5) is 3.2 x 10-9.
Thus the [H+] is 3.2 x 10-9 M
 Next calculate the [OH-]
 1.0 x 10-14 / 3.2 x 10-9 = 3.1 x 10-6 M
 Example 4: Calculate the pH of a solution of
household ammonia whose [OH-] is 7.93 x 10-3 M.
 Solution: This time you first calculate the [H+] from
the [OH-]
 7.93 x 10-3 M OH- = 1.26 x 10-12 M H+
 Then find the pH
 -log[1.26 x 10-12] = 11.9
 Now you try a few by yourself.
 Practice #1. What is the pH of a solution of NaOH that
has a [OH-] of 3.5 x 10-3 M?
 Practice #2. The H+ of vinegar that has a pH of 3.2 is
what?
 Practice #3. What is the pH of a 0.001 M HCl solution?
C.10.J distinguish between degrees of dissociation for
strong and weak acids and bases
Strong Acids:
 completely dissociate in water, forming H+
and an anion. example: HNO3 dissociates
completely in water to form H+ and N031-.
 The reaction is
 HNO3 → H+ + NO31 A 0.01 M solution of nitric acid contains 0.01
M of H+ and 0.01 M NO3- ions and almost no
HNO3 molecules. The pH of the solution
would be 2.0.
 There are only 6 strong acids: You must learn them.
The remainder of the acids therefore are considered
weak acids.
 HCl
 H2SO4
 HNO3
 HClO4
 HBr
 HI
 Note: when a strong acid dissociates only one H+ ion is
removed. H2SO4 dissociates giving H+ and HSO4- ions.
 H2SO4 → H+ + HSO41 A 0.01 M solution of sulfuric acid would contain 0.01 M
H+ and 0.01 M HSO41- (bisulfate or hydrogen sulfate
ion)
Weak acids:
 a weak acid only partially dissociates in water to give H+
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and the anion
for example, HF dissociates in water to give H+ and F-. It is a
weak acid. with a dissociation equation that is
HF ↔ H+ + FNote the use of the double arrow with the weak acid. That
is because an equilibrium exists between the dissociated
ions and the undissociated molecule. In the case of a
strong acid dissociating, only one arrow ( → ) is required
since the reaction goes virtually to completion.
An equilibrium expression can be written for this system:
Ka = [ H+][F-] / [HF]
 Which are the weak acids? Anything that
dissociates in water to produce H+ and is not one
of the 6 strong acids.
 Molecules containing an ionizable proton. (If the
formula starts with H then it is a prime candidate for
being an acid.) Also: organic acids have at least one
carboxyl group, -COOH, with the H being ionizable.
 Anions that contain an ionizable proton.
HSO41- → H+ + SO42 Cations: (transition metal cations and heavy metal
cations with high charge)
 also NH4+ dissociates into NH3 + H+
Strong Bases:
 They dissociate 100% into the cation and OH-
(hydroxide ion).
 example: NaOH → Na+ + OH a. 0.010 M NaOH solution will contain 0.010 M OHions (as well as 0.010 M Na+ ions) and have a pH of 12.
 Which are the strong bases?
 The hydroxides of Groups I and II.
 Note: the hydroxides of Group II metals produce 2
mol of OH- ions for every mole of base that
dissociates. These hydroxides are not very soluble, but
what amount that does dissolve completely dissociates
into ions.
exampIe: Ba(OH)2 → Ba2+ + 2OH a. 0.000100 M Ba(OH)2 solution will be 0.000200 M in
OH- ions (as well as 0.00100 M in Ba2+ ions) and will
have a pH of 10.3.
Weak Bases:
 What compounds are considered to be weak
bases?
 Most weak bases are anions of weak acids.
 Weak bases do not furnish OH- ions by dissociation.
They react with water to furnish the OH- ions.
 Note that like weak acids, this reaction is shown to be
at equilibrium, unlike the dissociation of a strong base
which is shown to go to completion.
 When a weak base reacts with water the OH- comes from
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the water and the remaining H+ attaches itself to the weak
base, giving a weak acid as one of the products. You may
think of it as a two-step reaction similar to the hydrolysis of
water by cations to give acid solutions. examples:
NH3 + H2O → NH4+ + OHmethylamine: CH3NH2+ H2O → CH3NH3+ + OHacetate ion: C2H3O2- + H2O → HC2H3O2 + OHGeneral reaction:
weak base + H2O → weak acid + OHSince the reaction does not go to completion relatively few
OH- ions are formed.