Acids and Bases – Unit 13
Chemistry of Acids and Bases
1. Watch video and complete worksheet

Standard Deviants Teaching Systems: Chemistry:
Module 05: Acids and Bases

http://app.discoveryeducation.com/player/view/asset
Guid/DBD191DB-A10E-43C2-8DDE-A73858F12FE2
2. Gallery walk to complete notes on pages
3-5 in packet
3. Homework is on page 6 in packet
Unit 13 – Acids and Bases
Notes #1: Intro
Acids: Something that produces a
hydrogen ion (H+) in solution
Unit 13 – Acids and Bases
Notes #1: Intro
• Properties of Acids:
• Tart or sour taste (lemon juice)
• Electrolytic
• Both strong and weak
• Will cause indicators to change colors
• A metal + an acid will produce hydrogen
gas
• Single replacement reaction
• Acid + metal → hydrogen gas + a “salt”
• Double replacement reaction
• Acid + Base → water + a “salt”
• Single replacement reaction
Acid + Metal → __Hydrogen gas_ + a “_salt_”
• Double replacement reaction
Acid + Base → _water__ + a “_salt_”
Remembering Acid Naming
Rules
“Handle acids carefully so you don’t get a case of “ate-icite-ous.””
 Polys
inENDING
“-ate” are changed
to “-ic”
ION
TYPEending
ION
ACID NAME
BEGINNING
Polys ending in-ite
“-ite” are charged
to “-ous”
NO hydrobeginning
Polyatomic
-ate
NO hydro- beginning

Monatomic
-ide
hydro- beginning
Hydro- prefix is not used
with poly containing
acids!!!!!
ACID ENDING
-ous
-ic
-ic
Examples of Naming Binary
Acids

HCl

HF
Hydrochloric acid
Hydrofluoric acid

HBr
Hydrobromic acid
Examples of Naming Ternary Acids

H2SO4

H2CO3

H2NO2
Sulfate is the poly, so sulfuric
carbonate is the poly, so
acid
carbonic acid
Nitrite is the poly, so nitrous
acid
• Base: Something that produces a hydroxide
ion (OH-) in solution
Unit 13 – Acids and Bases
• Properties of Bases:
• bitter
• slippery (soap)
• electrolytic
• Both strong and weak
• Will cause an indicator to
change colors
• Naming Bases
•
The easiest are the bases, since most of these are
_metal hydroxides, compounds you already know
how to name.
• Metal hydroxides are named in the same way any
other ionic compound is named. First give the name
of the _metal_ ion. Follow this with the name of the
anion, which, in the case of bases, is
“__hydroxide__”.
• KOH – Potassium Hydroxide
• Mg(OH)2 – Magnesium Hydroxide
Other definitions of Acids and
Bases

Arrhenius Acids and Bases:
 Acid:
 Hydrogen
containing compound that
ionize to yield a hydrogen ion in
solution.
 Base:
 Compounds
that ionize to yield a
hydroxide ion in solution.
Brønsted – Lowry Acids and
Bases

They felt the Arrhenius definition was too
limiting.
 Acids:
 Hydrogen
ion donor (Proton donor)
 Bases:
 Hydrogen
ion acceptor (Proton acceptor)
Brønsted – Lowry Acids and
Bases

Examples:
 NH3
+ H2O ↔ NH4+ + OH-

H2O donated the H+ - Acid

NH3 accepted the H+ - Base
 HCl
+ H2O ↔ H3O + + Cl-

HCl donated the H+ - Acid

H2O accepted the H+ - Base

Amphoteric:

Substance that can act as both an acid or a base.

Background Theory:

The oxides of metals are basic in nature. For example, the oxides of
the alkali metals (Group I) form alkali or basic solutions.
o
Sodium oxide + water → Sodium hydroxide solution
Na2O(s) + H2O(l) → NaOH(aq)

The soluble oxides of non-metals are acidic in nature. Examples
include, carbon dioxide, sulfur dioxide and nitrogen dioxide.
o
Sulfur dioxide + water → Sulfurous acid
SO2(g) + H2O(l) → H2SO3(aq)
o
Insoluble non-metallic oxides like carbon monoxide do not form
acidic solutions. This is often the cause of acid rain.
Compounds such as the amino acids, which contain both acidic and basic
groups in their molecules, can also be described as amphoteric.
Strong Acids and Bases

Strong Acids/Bases:

Those that ionize completely in solution.


Ex: HCl, NaOH
Weak Acids/Bases:

Those that only slightly ionize in solution.

Ex: NH3, Acetic Acid (vinegar)

Tooth decay is caused by the weak acid – lactic
acid: C3H6O3
Homework: pg 6
The pH Scale
pg 7-
pH Scale
MEASURING pH
Scientists use a pH scale to measure the strength of an acid or base. The
term pH stands for “potential for hydrogen”. The amount of hydrogen
in a substance determines its acidity or alkalinity. Alkaline is another term
for base. A number on the pH scale is used to describe the strength of
acidity or alkalinity. The most commonly used pH scale goes from 1 (very
acidic) to 14 ( very basic). The number 7 on a pH scale means neutral –
neither acid nor base.
Acids play important roles in the chemistry of living things. Many of the
foods you eat are acids in vitamins like ascorbic acid or vitamin C, and
folic acid. Other acids help the body such as stomach acids and others
are waste products of cell processes like lactic acid in working muscles.
Acids also are used to make valuable products for homes, farms and
industries. People often use dilute solutions of acids to clean brick and
other surfaces. Hardware stores sell muriatic (hydrochloric ) acid, which
is used to clean bricks and metals. Industry uses sulfuric acid in car
batteries, to refine petroleum and to treat iron and steel. Farmers
depend on the nitric acid and phosphoric acid to make fertilizers for
crops, lawns, and gardens.
The concentration of hydrogen ions in a solution
is described by its number on the pH scale.
• A low pH tells you that the concentration
of hydrogen ion is high.
EX: pH 2
• By comparison, a high pH tells you that the
concentration of hydrogen ion is low.
EX: pH 12
Self-ionization of water
 Self-ionization
of water:
Reaction
in which 2 water
molecules produce ions
H2O
+ H2O → OH- + H3O+
Also
written as: H2O ↔ H+ + OH-
The
H3O+ and H+ represent hydrogen
ions in solution.
Neutral Solutions
 In
pure water, the concentration of
hydrogen ions is equal to the
concentration of hydroxide ions
1
x 10-7M or pH of 7
Remember
 [H+]
= [OH-]
(brackets
 This
M represents Molarity
represent concentration)
represents a neutral solution.
Solutions
In a solution, if the [H+] increases, the [OH-] decreases and
vice versa.
 Think back to a see-saw. As one person went up the other
went down.
 Ion-product constant of water, Kw:

Kw


= [H+] x [OH-] = 1 x 10-14M
Acidic Solution:
 The [H+] will be greater than the [OH-].
 Therefore, the [H+] is greater than 1 x 10-7M.
 Think about the # line. -5 is GREATER than -7
Basic Solution:
 The [H+] will be less than [OH-].
 Therefore, the [H+] is less than 1 x 10-7M.
 A.k.a. alkaline solutions
NUMBER LINE and pH
 Remember
the number line
Increasing
-7 -6 -5 -4 -3 -2 -1 0 1 2 3 4 5 6 7 8

Which is greater? 0 or 3


Which is greater? -7 or -4


3
-4
Which is less? -2 or -4

-4
Acids
Bases
Homework pg. 9
pH Calculations

The pH scale ranges from 0-14.

0 = very acidic

7 = neutral

14 = very basic

pH = -log [H+]

What is the pH of a neutral solution?

Calculate using the Logarithmic function on
the calculator (see at right)
Sample Problems

As long as you have a 1 x 10 to some power, the pH is
the exponent.
1. What is the pH of the following concentrations?
a. [H+] = 1 x 10-2M
pH = 2 acidic
b. [H+] = 1 x 10-9M
pH = 9 basic
c. [H+] = 1 x 10-5M
pH = 5 acidic
Sample Problems

If you do not have 1 to the power then you MUST use
our formulas.
2. What is the pH of the following?
a. [H+] = 2x10-2

pH = -log(2x10-2) = 1.7 pH
b. [H+] = 6x10-9

pH = -log(6x10-9) = 8.2 pH
c. [H+] = 3x10-5

pH = -log(3x10-5) = 4.5 pH
Other Formulas and Problems

pH 14 = pH + pOH
(See example 1 in Example Problems))

Equilibrium constant labeled as Kw

Kw is 1x10-14

Kw = [OH-] x [H+] = 1x10-14
Other Formulas and Problems
EX: What is the pH of a solution with a
4.0 x 10-11M?
o
[OH-] of
Use Kw to find [H+] then find pH using –log function.
Step1:
Kw = [OH-] x [H+] = 1x10-14
[H+] = 1x10-14/4x10-11 = 2.5x10-4
Step 2:
pH = -log [H+]
pH= -log(2.5x10-4) = 3.6
1. If pH = 5, pOH =
pH 14 = pH + pOH
14 = 5 + pOH
14 – 5 = 9 pOH
Acid because pH = 5
2. What is the pH of a solution that has a
hydrogen ion concentration of 1.0 x 10-5M?
Is this solution acidic, basic or neutral?
Given: [H+]
Solving for: pH
pH = - log [H+]
pH = - log(1.0 x 10-5 M)
pH = 5
pH < 7
ACIDIC
3. What is the hydrogen ion concentration of a
solution with a pH of 11? Which has a greater
concentration: H+ or OH-?
[H+] = 1 x 10 -11 M
more OH-, So basic
4. What is the pH of a solution that
has a hydrogen ion concentration of
1.2 x 10-8M? Is this solution acidic,
basic or neutral?
Given: [H+] Solving for: pH
pH = - log [H+]
pH = - log(1.2 x 10-8 M)
pH = 7.92
pH > 7
BASIC
5. Assuming Kw = 1x10-14, calculate the
molarity of OH- in solutions at 25ºC when the
H+ concentration is 0.2M
At 25ºC,
Kw = [OH-] [H+] = 1x10-14
1x10-14 = [OH-] 0.2M
= 1x10-14/ .2
[OH-] = 5x10-14 M
HOMEWORK: pg 12
Neutralization Notes

Acid-Base reactions will produce salt water
when completely neutralized.

Salts are compounds consisting of a(n) anion
from an acid and a(n) cation from a base.

In general, reactions in which an acid and a
base react in an aqueous solution to produce a
salt and water is called Neutralization
Reactions.
Neutralization Reactions

Neutralization occurs when an
Acid + Base ↔ water + salt
 Salt:
Anion from acid and the cation from the base
join together to form a salt.
Where do we see this process?
• Antacids
• Farmers controlling the pH of soil
• Formation of caves
 A strong acid + a strong base = neutral solution
Examples:
HCl + NaOH ↔ H2O + NaCl
HCl + KOH ↔ H2O + KCl
Practice: Don’t forget to balance
them after you write them.

HCl + LiOH →

HNO3 + CsOH →

HBr + KOH →
HOH + LiCl
CsNO3 + H2O
H2O + KBr
Titrations

Titration: The process of adding a known amount
of solution of known concentration to determine
the concentration of the other solution.

If you don’t know the concentration of one
solution, you can figure it out by performing a
neutralization reaction, or titration, with a
standard solution.

A standard solution is one of known
concentration.
Performing Titrations

Steps in a neutralization reaction:
 1.
A measured volume of an acid
solution of unknown concentration is
added to a flask.
 2.
Several drops of indicator are
added to the solution.
 3.
Measured volumes of a base with a
known concentration are mixed into
the acid until it barely changes color.
Performing Titrations, cont.

End Point: The point at which the
indicator changes color.

Once you have reached the end point, you can
perform calculations to find the unknown
solution.
Let’s show a video!
http://app.discoveryeducation.com/search?
Ntt=titration##
http://www.youtube.com/watch?v=RHTxIY
DJ730
Performing Titrations, cont.

Example: A 25 mL solution of H2SO4 is completely neutralized by 18
mL of 1.0 M NaOH. What is the concentration of H2SO4 solution?
 Step 1: Balanced equation


2
2 2O
____H2SO4 + ____NaOH
↔ ____Na2SO4 + ____H
Step 2: Use formula to solve for unknown.
 MaVa = MbVb
na
 na =
nb
Number of moles of your Acid (coefficient)
 nb = Number of moles of your Base (coefficient)
 M = Molarity of acid or base
 V = Volume of acid or base (in Liters)
MaVa = MbVb
na
nb
Ma ( 25 mL)
1 mol
=
(1.0 M)( 18 mL)
2 mol
Molarity = 0.36 M
1. How many moles of HCl are needed to
neutralize 6 mols of KOH?

1st ask, what is the mol ratio and then set it up as a
proportion.
HCl + KOH  KCl + H2O
This equation is balanced so
1 mole HCl = 1 mole KOH
So 6 mols KOH will neutralize
6 moles HCl
2.
H2SO4 + 2NaOH ↔ Na2SO4 + 2H2O
2 moles of NaOH.
a. One mole of sulfuric acid is needed to neutralize
b. How many moles of NaOH are needed to neutralize 4 moles of H2SO4?
Given that 1 H2SO4 = 2NaOH
So if you have 4 mols H2SO4
you will need 8
moles NaOH
Homework pg 17
Also begin working on your
Review on pages: 18-21