Chemistry
Chemistry is the study of matter—its composition,
properties, and reactions.
Matter is anything that has mass and takes up volume.
Naturally occurring matter:
• cotton
• sand
• digoxin, a cardiac drug
Synthetic (human-made) matter:
• nylon
• Styrofoam
• ibuprofen
1
States of Matter.
Three States of Matter—Solid, Liquid, and Gas
Solid:
•
has a definite volume,
•
maintains its shape regardless of its container, and
•
has particles that lie close together in a regular
three-dimensional array.
•
Mostly high density
2
Liquid:
• has definite volume,
• takes the shape of its container, and
• has particles that are close together, but they can
randomly move around, sliding past one another
• Relatively low density
3
Gas:
• has no definite shape or volume,
• expands to fill the volume and assumes the
shape of whatever container it is put in, and
• has particles that are very far apart and move
around randomly.
• Very low density
4
Properties of Matter
Physical properties can be observed or measured
without changing the composition of the material.
• boiling point (bp)
• melting point (mp)
• solubility
• color
• odor
5
Physical change alters the material without changing
its composition:
Change of state or density may occur. Substance retains
its properties.
• melting ice (solid water) to form liquid water
• boiling liquid water to form steam (gaseous water).
6
Chemical properties determine how a substance can
be converted into another substance.
Chemical change, or chemical reaction, converts
one substance into another.
Formed substance/s has different properties from initial
substances.
For example:
• a piece of paper burning,
• metabolizing an apple for energy, or
• oxygen and hydrogen combining to form water.
7
• Determine whether physical or chemical change
forming ice cubes
burning natural gas
snow melting
fermenting grapes to form wine
8
Classification of Matter
Pure Substance and Mixtures
•Pure Substance
• is composed of a single component,
• has a constant composition,
• cannot be broken down to other pure substances
by a physical change.
• Table sugar (C12H22O11) and water (H2O) are both
pure substances.
9
Mixture
• is composed of more than one substance
• can have varying composition (any combination
of solid, liquid, and gas), depending on the sample
• can be separated into its components by a
physical change
• sugar dissolved in water = mixture
10
Element vs. Compound
An element is a
pure substance that
cannot be broken down
by a chemical change.
aluminum metal (Al)
A compound is a
pure substance formed
by chemically joining
two or more elements.
table salt (NaCl)
11
Classification of Matter
12
Measurement
Every measurement is composed of a number and a unit.
13
The Importance of Units
A number is meaningless without the unit.
• proper aspirin dosage = 325 (milligrams or kilograms?)
• a fast time for the 100-meter dash = 10.00 (seconds
or days?)
The English system uses units like miles (length),
gallons (volume), and pounds (weight).
The metric system uses units like meters (length),
liters (volume), and grams (mass).
14
The Metric System of Units
Each type of measurement has a base unit.
15
• Other units are related to the base unit by a power of 10.
• The prefix of the unit name indicates if the unit is larger
or smaller than the base unit.
16
CONVERSIONS.
micro (µm) = (10x-6)m
x
mm
cm
dm
m
Dm
Hm
Km
/
mega (Mm) = (10x6)m
e.g. 1 cm = ? mm
= 10 mm (1 centimeter = 10 millimeters)
1 m = ? mm
= 1000 mm (1 meter = 1000 millimeters)
1 m = ? Km
= 1/1000 Km
= 0.001 Km (1 meter = 0.001 kilometer
Measuring Mass
Mass is a measure of the amount of matter in an
object.
Weight is the force that matter feels due to gravity.
1,000 grams (g) = 1 kilogram (kg)
1 g = 0.001 kg
1 gram (g) = 1,000 milligram (mg)
0.001 g = 1 mg
18
Measuring Volume
1,000 liters (L) = 1 kiloliter (kL)
1 L = 0.001 kL
1 liters (L) = 1,000 milliliter (mL)
0.001 L = 1 mL
Volume = Length x Width x Height
= cm x cm x cm
= cm3
1 mL = 1 cm3 = 1 cc
19
Significant Figures
Exact and Inexact Numbers
An exact number results from counting objects or is
part of a definition.
• 10 fingers
• 10 toes
• 1 meter = 100 centimeters
An inexact number results from a measurement or
observation and contains some uncertainty.
• 15.3 cm
• 1000.8 g
• 0.0034 mL
20
Significant Figures
Determining Significant Figures
Significant figures are all the digits in a measured
number including one estimated digit.
All nonzero digits are always significant.
65.2 g
3 sig. figures
255.345 g
6 sig. figures
21
Significant Figures
Rules to Determine When a Zero
is a Significant Figure
Rule 1: A zero counts as a significant figure when
it occurs:
• between two nonzero digits, as in
29.05 g
4 sig. figures
1.0087 mL
5 sig. figures
• or at the end of a number with a decimal place.
3.7500 cm
5 sig. figures
620. lb
3 sig. figures
22
Significant Figures
Rules to Determine When a Zero
is a Significant Figure
Rule 2: A zero does not count as a significant figure
when it occurs:
• at the beginning of a number,
0.00245 mg
3 sig. figures
0.008 mL
1 sig. figure
• at the end of a number that does not have a decimal.
2570 m
3 sig. figures
1245500 m
5 sig. figures
23
Significant Figures
Rules for Multiplication and Division
The answer has the same number of significant figures
as the original number with the fewest significant
figures.
4 sig. figures
351.2 miles
=
5.5 hour
2 sig. figures
63.854545 miles
hour
The answer must have
2 sig. figures.
24
Significant Figures
Rules for Rounding Off Numbers
to be retained
to be dropped
63.854545 miles =
hour
first digit to be dropped
If the first digit
to be dropped is:
64 miles
hour
2 sig. figures
Answer
Then:
• between 0 and 4
• drop it and all remaining digits
• between 5 and 9
• round up the last digit
to be retained by adding 1
25
Significant Figures
Rules for Multiplication and Division
23.2
x
3 sig. figs
25.0
3 sig. figs
1.1
=
2 sig. figs
x
0.50
=
2 sig. figs
25.52
26
calculator
display
2 sig. figs
50
50.
calculator
display
2 sig. figs
26
Significant Figures
Rules for Addition and Subtraction
The answer has the same number of decimal places
as the original number with the fewest decimal places.
10.11 kg 2 decimal places
=
3.6 kg
1 decimal place
6.51 kg
answer must have
1 decimal place
6.5 kg
final answer
1 decimal place
27
Scientific Notation
In scientific notation, a number is written as:
y x 10x
Coefficient:
A number between
1 and 10
Exponent:
Any positive
or negative
whole number
28
Scientific Notation
HOW TO Convert a Standard Number to Scientific Notation
Example
Convert these numbers to scientific notation.
2,500
Step [1]
Move the decimal point to give a number
between 1 and 10.
2500
Step [2]
0.036
0.036
Multiply the result by 10x, where
x = number of places the decimal was moved.
• move decimal left,
x is positive
2.5 x 103
• move decimal right,
x is negative
3.6 x 10−2
29
Scientific Notation
Converting a Number in Scientific
Notation to a Standard Number
• When the exponent x is positive, move the
decimal point x places to the right.
2.800 x 102 = 280.0
• When the exponent x is negative, move the
decimal point x places to the left.
2.80 x 10–2 = 0.0280
30
Conversion Factors
Conversion factor is a term that converts a quantity in
one unit to a quantity in another unit.
original
x
quantity
conversion factor
desired
= quantity
• Conversion factors written as equalities.
2.20 lb = 1 kg
• To use them, they must be written as fractions.
2.20 lb
1 kg
or
1 kg
2.20 lb
31
Solving a Problem Using One Conversion Factor
Factor-label method uses conversion factors to
convert a quantity in one unit to a quantity in
another unit:
• units are treated like numbers.
• make sure all unwanted units cancel.
32
How many grams of aspirin are in a 325-mg
Example
tablet?
Step [1]
Identify the original quantity and the desired
quantity, including units.
original quantity
325 mg
desired quantity
?g
33
Step [2] Write out the equality needed to solve
the problem.
1 g = 1000 mg
This can be written as two possible factors:
1000 mg
1g
or
1g
1000 mg
Choose this factor to
cancel the unwanted
unit, mg.
34
Step [3]
Set up and solve the problem.
325 mg x
3 sig. figures
Step [4]
1g
1000 mg
=
Unwanted unit
cancels
0.325 g
3 sig. figures
Write the answer with the correct number
of significant figures.
35
Problem Solving Using Clinical
Conversion Factors
Sometimes conversion factors do not have to be
looked up or defined; they are stated in the problem.
• A patient is prescribed 1.25 g of amoxicillin, which is
available in 250-mg tablets. How many tablets are
needed?
1.25 g
x
1000 mg x 1 tablet = 5 tablets
250 mg
1g
36
Temperature
Temperature is a measure of how hot or cold an object is.
Three temperature scales are used:
• degrees Fahrenheit (°F)
• degrees Celsius (°C)
• Kelvin (K)
To convert from °C to °F:
TF = 1.8(TC) + 32
To convert from °F to °C:
TC = TF − 32
1.8
To convert from °C to K:
To convert from K to °C:
TK = TC + 273
TC = TK − 273
37
Temperature
Comparing the Three Temperature Scales
38
Density and Specific Gravity
Density
Density is a physical property that relates the mass of
a substance to its volume.
density
=
To convert volume (mL)
to mass (g):
g
mL x
= g
mL
density
mass (g)
volume (mL or cc)
To convert mass (g)
to volume (mL):
mL
g x
= mL
g
inverse of density
39
Density and Specific Gravity
Solving Problems with Density
• Calculate the mass in grams of 15.0 mL of a saline
solution that has a density 1.05 g/mL.
15.0 mL
original quantity
?g
desired quantity
• Density is the conversion factor, and can be
written two ways:
1.05 g
1 mL
1 mL
1.05 g
Choose the inverse density
to cancel the unwanted unit, mL.
40
Density and Specific Gravity
Solving Problems with Density
• Set up and solve the problem:
15.0 mL
x
1.05 g
1 mL
=
3 sig. figures
15.8 mL
3 sig. figures
Unwanted unit
cancels
• Write the final answer with the correct number
of significant figures.
41
Density and Specific Gravity
Specific Gravity
Specific gravity is a quantity that compares the density
of a substance with the density of water at the
same temperature.
specific gravity
=
density of a substance (g/mL)
density of water (g/mL)
• The units of the numerator (g/mL) cancel the
units of the denominator (g/mL).
• The specific gravity of a substance is equal to its
density, but contains no units.
42
Atoms and the Periodic Table
Elements
• An element is a pure substance that cannot be
broken down into simpler substances by a
chemical reaction.
• Each element is identified by a one- or two-letter
symbol.
• Elements are arranged in the periodic table.
• The position of an element in the periodic table
tells us much about its chemical properties.
43
Atoms and the Periodic Table
Elements
44
Atoms and the Periodic Table
45
Atoms and the Periodic Table
Metals, Nonmetals, and Metalloids
The elements in the periodic table are divided into
three groups—metals, nonmetals, and metalloids.
Metals:
• are located on the left side of the periodic table,
• usually exist as shiny solids,
• are good conductors of heat and electricity, and
• are solids at room temperature, except for
mercury (Hg), which is a liquid.
46
Atoms and the Periodic Table
Metals, Nonmetals, and Metalloids
Nonmetals:
• are located on the right side of the periodic table,
• usually do not have a shiny appearance,
• are usually poor conductors of heat and electricity, and
• can be solids, liquids, or gases at room temperature.
solid
sulfur
carbon
liquid
bromine
gas
nitrogen
oxygen
47
Atoms and the Periodic Table
Metals, Nonmetals, and Metalloids
Metalloids:
• are located on the solid line that starts at boron
(B) and angles down towards astatine (At),
• have properties intermediate between metals
and nonmetals,
• are represented by only seven elements.
boron (B)
silicon (Si)
germanium (Ge)
arsenic (As)
antimony (Sb)
tellurium (Te)
astatine (At)
48
Focus on the Human Body
The Elements of Life
Four nonmetals, O, C, H, and N, comprise 96% of the
mass of the human body and are called the buildingblock elements.
49
Atoms and the Periodic Table
Compounds
Compound: a pure substance formed by chemically
combining two or more elements together.
A chemical formula consists of:
• element symbols to show the identity of the
elements forming a compound, and
• subscripts to show the ratio of atoms in the
compound.
H2O
2 H atoms 1 O atom
C3H8
3 C atoms
8 H atoms
50
Atoms and the Periodic Table
Compounds
• Compounds can be drawn many ways:
• Different elements are represented by different
colors:
51
Structure of the Atom
All matter is composed of the same basic building
blocks called atoms.
Atoms are composed of three subatomic particles:
52
Structure of the Atom
Nucleus:
• location of protons
and neutrons
• dense core of the atom
• most of the mass of the
atom resides here
Electron cloud:
• location of electrons
• comprises most of the
atom’s volume
53
Structure of the Atom
Atomic Number
From the periodic table:
3
Li
Atomic number (Z) is
the number of protons
in the nucleus.
• Every atom of a given element has the same
atomic number.
• Every atom of a given element has the same
number of protons in the nucleus.
• Different elements have different atomic numbers.
• A neutral atom has no net overall charge, so
Z = number of protons = number of electrons.
54
Mass Number
the number of protons (Z)
Mass number (A) =
+
the number of neutrons
55
Isotopes
Isotopes, Atomic Number, and Mass Number
Isotopes are atoms of the same element that have
a different number of neutrons.
Mass number (A)
35
Atomic number (Z)
17
# of protons
=
# of electrons =
# of neutrons =
=
Cl
37
Cl
17
17
17
17
35 – 17
18
17
37 – 17
20
56
Isotopes
Atomic Weight
The atomic weight is the weighted average of the
masses of the naturally occurring isotopes of a
particular element reported in atomic mass units.
From the periodic table:
82
Pb
207.2
atomic number (Z)
element symbol
atomic weight (amu)
57
The Periodic Table
Basic Features of the Periodic Table
A row in the periodic table is called a period, and a
column in the periodic table is called a group.
Main group elements:
• consist of the two columns on the far left and the
six columns on the far right of the periodic table;
• the groups are numbered 1A–8A.
Transition metal elements:
• contained in the 10 short columns in the middle;
• these groups are numbered 1B–8B.
Inner transition elements:
• consist of the lanthanides and actinides;
• no group numbers are assigned.
58
The Periodic Table
Basic Features of the Periodic Table
59
The Periodic Table
Characteristics of Groups 1A and 2A
Elements that comprise a particular group have
similar chemical properties.
Group
Number
Group
Name
1A
Alkali metals
2A
Alkaline earth
elements
Properties of Both Groups
• soft and shiny metals
• low melting points
• good conductors of heat
and electricity
• react with water to form
basic solutions
60
The Periodic Table
Characteristics of Groups 7A and 8A
Group
Number
Group
Name
7A
Halogens
8A
Noble gases
Properties
• exist as two atoms
joined together
• very reactive; combine
with many other elements
to form compounds
• very stable
• rarely combine with
any other elements
61
The Unusual Nature of Carbon
Carbon is different from most other elements in that it
has three elemental forms:
• Diamond consists
of a 3-dimensional
network of C atoms.
• Graphite
contains parallel
sheets of C
atoms.
• Buckminsterfullerene
contains a sphere with
60 C atoms.
62
Electronic Structure
The chemistry of an element is determined by the
number of electrons in an atom.
• Electrons do not move freely in space; rather they
occupy specific energy levels.
• The regions occupied by electrons are called
principal energy levels or shells (n).
• The shells are numbered n = 1, 2, 3, etc.
• Electrons in lower numbered shells are closer to
the nucleus and are lower in energy.
• Electrons in higher numbered shells are further
from the nucleus and are higher in energy.
63
Electronic Structure
Shells
Shells with larger numbers (n) are farther from the
nucleus, have a larger volume, and therefore, can
hold more electrons.
The distribution of electrons in the first four shells:
increasing
energy
Shell (n)
Number of Electrons
in a Shell
4
32
3
18
2
8
1
2
increasing
number of
electrons
64
Electronic Structure
Subshells and Orbitals
• Shells are divided into subshells, identified by the
letters s, p, d, and f.
• The subshells consist of orbitals.
• An orbital is a region of space where the
probability of finding an electron is high.
• Each orbital can hold two electrons.
increasing
energy
Subshell
s
p
Number of Orbitals
1
3
d
5
f
7
65
Electronic Structure
Subshells and Orbitals
66
Electronic Structure
Orbital Shapes
• The s orbital has a spherical shape.
• The p orbital has a dumbbell shape.
67
Electron Configuration
The electron configuration describes how the electrons
are arranged in an atom’s orbitals. The lowest energy
arrangement is called the ground state.
Rules to Determine the Ground State Electronic
Configuration of an Atom
Rule [1] Electrons are placed in the lowest energy
orbitals beginning with the 1s orbital.
• An orbital closer to the nucleus is lower in energy.
• Within a shell, orbital energies increase in the
following order: s, p, d, f.
• These guidelines result in the following order of
energies 1s, 2s, 2p, 3s, 3p.
Rule [2] Each orbital holds a maximum of 2 electrons. 68
Electron Configuration
Rules to Determine the Ground State Electronic
Configuration of an Atom
69
70
Electron Configuration
First-Row Elements (Period 1)
Element
Electron
Configuration
H (Z = 1)
1 electron
1s1
He (Z = 2)
2 electrons
1s2
71
Electron Configuration
Second-Row Elements (Period 2)
Element
Li (Z = 3)
3 electrons
Electron
Configuration
1s22s1
C (Z = 6)
6 electrons
1s22s22p2
Ne (Z = 10)
10 electrons
1s22s22p6
72
Valence Electrons
• The chemical properties of an element depend on
the most loosely held electrons, which are found
in the valence shell.
• The valence shell is the outermost shell (the highest
value of n).
• The electrons in the valence shell are called valence
electrons.
Be
Cl
1s22s2
1s22s22p63s23p5
valence shell: n = 2
valence shell: n = 3
# of
valence electrons = 2
# of
valence electrons = 7 73
Valence Electrons
• Elements in the same group have the same
number of valence electrons.
• The group number, 1A–8A, equals the number of
valence electrons for the main group elements.
• The exception is He, which has only 2 valence
electrons.
• The chemical properties of a group are similar
because these elements contain the same
electronic configuration of valence electrons.
74
Valence Electrons
Main Group Elements
Group number:
1A
2A
3A
Period 1:
H
1s1
4A
5A
6A
7A
8A
He
1s2
Period 2:
Li
2s1
Be
B
C
N
O
F
Ne
2s2 2s22p1 2s22p2 2s22p3 2s22p4 2s22p5 2s22p6
Period 3:
Na
Mg
Al
Si
P
S
Cl
Ar
3s1 3s2 3s23p1 3s23p2 3s23p3 3s23p4 3s23p5 3s23p6
75
Valence Electrons
Electron-Dot Symbols
• Dots representing valence electrons are placed
on the four sides of an element symbol.
• Each dot represents one valence electron.
• For 1–4 valence electrons, single dots are used.
With > 4 valence electrons, the dots are paired.
Element:
H
C
O
Cl
# of Valence electrons:
1
4
6
7
Electron-dot symbol:
H
C
O
Cl
76
Periodic Trends
Atomic Size
Increases
• The size of atoms
increases down a
column, as the
valence e− are
farther from the
nucleus.
Decreases
• The size of atoms decreases across a row, as
the number of protons in the nucleus increases.
• The increasing # of protons pulls the e− closer
to the nucleus, making the atoms smaller.
77
Periodic Trends
Ionization Energy
The ionization energy is the energy needed to remove
an electron from a neutral atom.
Na + energy  Na+ + e–
Decreases
Increases
• Ionization energies
decrease down a
group as the
valence e− get
farther away from
the positively
charged nucleus.
• Ionization energies increase across a period as the
number of protons in the nucleus increases.
78
Introduction to Bonding
Ionic and Covalent Compounds
Bonding is the joining of two atoms in a stable
arrangement.
Only the noble gases in group 8A do not readily
react to form bonds, because their electronic
configuration is especially stable.
Elements will gain, lose, or share electrons to attain
the electronic configuration of the noble gas closest
to them in the periodic table.
79
Ionic and Covalent Compounds
Introduction to Bonding
There are two different kinds of bonding:
• Ionic bonds result from the transfer of electrons
from one element to another.
• Covalent bonds result from the sharing of
electrons between two atoms.
The position of an element in the periodic table
determines the type of bonds it makes.
80
Ionic Compounds
Ionic Bonding
Ionic bonds form between:
• a metal on the left side of the periodic table and
• a nonmetal on the right side of the periodic table.
Na
sodium
metal
+
Cl2
chlorine
gas
NaCl
sodium chloride
crystals
81
Covalent Bonding
Covalent bonds are formed when two nonmetals
combine, or when a metalloid bonds to a nonmetal.
A molecule is a discrete group of atoms that share
electrons.
82
Ions
Cations and Anions
• Ions are charged species in which the number
of protons and electrons in an atom is unequal.
• Ionic compounds consist of oppositely charged
ions that have a strong electrostatic attraction
for each other.
• There are two types of ions—cations and anions.
83
Ions
Cations and Anions
Cations are positively charged ions. A cation has
fewer electrons (e−) than protons. Metals form cations.
84
Ions
Cations and Anions
Anions are negatively charged ions. An anion has
more e− than protons. Nonmetals form anions.
85
Ions
The Octet Rule
By losing one, two, or three e−, an atom forms a cation
with a completely filled outer shell of e−.
By gaining one, two, or three electrons, an atom forms
an anion with a completely filled outer shell of e−.
The octet rule: a main group element is especially
stable when it possesses an octet of e− in its outer
shell.
octet = 8 valence e−
86
Ions
Relating Group Number to Ionic Charge for
Main Group Elements
• Elements in the same group form ions of similar
charge.
• Metals form cations.
• For metals in groups 1A, 2A, and 3A, the group
number = the charge on the cation.
• Nonmetals form anions.
• For nonmetals in Groups 5A, 6A and 7A, the anion
charge = the group number –8.
87
Ions
Relating Group Number to Ionic Charge for
Groups 1A–3A
the cation charge = the group number
group 1A:
M
1 valence e−
M + + e−
group 2A:
M
2 valence e−
M2+ + 2e−
group 3A:
M
3 valence e−
M3+ + 3e−
88
Ions
Relating Group Number to Ionic Charge for
Groups 5A, 6A, and 7A
the anion charge = group number – 8
group 5A:
group 6A:
group 7A:
X
+ 3e−
5 valence e−
X
+ 2e−
6 valence e−
X
+ e−
7 valence e−
X
3−
charge = 5 – 8 = – 3
X
2−
charge = 6 – 8 = – 2
X
−
charge = 7 – 8 = – 1
89
Ions
Ions Formed by the Main Group Elements
90
Ions
Metals with Variable Charge
91
Ionic Compounds
Formulas for Ionic Compounds
When a metal (on the left side of the periodic table)
transfers one or more electrons to a nonmetal
(on the right side), ionic bonds are formed.
The sum of the charges in an ionic compound must
be zero overall.
92
Ionic Compounds
HOW TO Write a Formula for an Ionic Compound
Step [3]
To write the formula, place the cation
first and then the anion, and omit charges.
CaO
CaCl2
• Use subscripts to show the number of
each ion needed to have a zero overall
charge.
93
Naming Ionic Compounds
Naming Cations
Main group cations are named for the element from
which they are formed.
Na+
sodium
K+
potassium
Ca2+
calcium
Mg2+
magnesium
94
Naming Ionic Compounds
95
Naming Ionic Compounds
Naming Anions
Anions are named by replacing the ending of the
element name by the suffix “-ide”.
96
Naming Ionic Compounds
Compounds of Main Group Metals
• Name the cation and then the anion.
• Do not specify the charge on the cation.
• Do not specify how many ions of each type are
needed to balance charge.
Na+
+
sodium
Mg2+
+
magnesium
F−
fluoride
Cl−
chloride
NaF
sodium fluoride
MgCl2
magnesium chloride
97
Naming Ionic Compounds
Writing a Formula from the Name
HOW TO Derive a Formula from the Name of an Ionic
Compound
Example
Write the formula for tin(IV) oxide.
Step [1]
Identify the cation and anion and
determine their charges.
tin(IV) oxide
Sn4+
O2−
98
Naming Ionic Compounds
Writing a Formula from the Name
HOW TO Derive a Formula from the Name of an Ionic
Compound
Step [2]
Balance charges.
Sn4+
Step [3]
O2−
Two −2 anions
are needed for
each +4 cation.
Write the formula with the cation first, and
use subscripts to show how many of each
ion is needed to have zero overall charge.
final answer =
SnO2
99
Physical Properties of Ionic Compounds
• Ionic compounds are crystalline solids that have
very high melting points (NaCl = 801 oC) and
extremely high boiling points (NaCl = 1413 oC).
• When ionic compounds dissolve in water, they
separate into cations and anions. The resulting
aqueous solutions conduct an electric current.

+
NaCl
water
solution
100
Polyatomic Ions
A polyatomic ion is a cation or anion that contains
more than one atom.
101
Polyatomic Ions
Writing Formulas for Ionic Compounds with
Polyatomic Ions
• When a cation and anion of equal charge
combine, only one of each ion is needed.
Na+ + NO2−
NaNO2
zero overall
charge
Ba2+ + SO42−
BaSO4
zero overall
charge
102
Polyatomic Ions
Writing Formulas for Ionic Compounds with
Polyatomic Ions
• When a cation and anion of unequal charge
combine, use the ionic charges to determine
the relative number of each ion that is needed.
Mg2+
+
OH−
+2 charge means −1 charge means
1 Mg2+ anion is
2 OH− anions are
needed
needed
Mg(OH)2
zero overall
charge
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Polyatomic Ions
Naming Ionic Compounds with Polyatomic Ions
The same rules are followed for naming standard
ionic compounds:
• Name the cation and then the anion.
• Do not specify the charge on the cation.
• Do not specify how many ions of each type are
needed to balance charge.
Na2CO3
sodium carbonate
BaSO4
barium sulfate
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Focus on Health & Medicine
Useful Ionic Compounds
Ionic compounds are the active ingredients in several over-the-counter
drugs: calcium carbonate (CaCO3), the antacid in Tums; magnesium
hydroxide [Mg(OH)2], an active component in Maalox; and iron(II)
sulfate (FeSO4), an iron supplement used to treat anemia. 105
Covalent Compounds
Introduction to Covalent Bonding
Covalent bonds result from the sharing of electrons
between two atoms.
• A covalent bond is a two-electron bond in which
the bonding atoms share the electrons.
• A molecule is a discrete group of atoms held
together by covalent bonds.
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Covalent Compounds
Introduction to Covalent Bonding
Unshared electron pairs are called nonbonded
electron pairs or lone pairs.
Atoms share electrons to attain the electronic
configuration of the noble gas closest to them
in the periodic table.
• H shares 2 e−.
• Other main group elements are stable when they
possess an octet of e− in their outer shell.
107
Covalent Compounds
Covalent Bonding and the Periodic Table
How many covalent bonds will a particular atom form?
• Atoms with one, two, or three valence e−
generally form one, two or three bonds,
respectively.
• Atoms with four or more valence electrons form
enough bonds to give an octet.
predicted
number of bonds
=
8 – number of valence e−
108
Covalent Compounds
Covalent Bonding and the Periodic Table
For most common atoms, except H
Number of bonds
+
Number of lone pairs
=
4
109
Naming Covalent Compounds
HOW TO Name a Covalent Molecule
Example
Name each covalent molecule:
(a) NO2
Step [1]
(b) N2O4
Name the first nonmetal by its element
name and the second using the suffix
“-ide”.
(a) NO2
nitrogen oxide
(b) N2O4
nitrogen oxide
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Naming Covalent Compounds
HOW TO Name a Covalent Molecule
Step [2]
Add prefixes to show the number of
atoms of each element.
• The prefix “mono-” is usually omitted.
Exception: CO is named carbon monoxide
mono + oxide = monoxide
111
Naming Covalent Compounds
HOW TO Name a Covalent Molecule
(a) NO2
nitrogen dioxide
(b) N2O4
dinitrogen tetroxide
112
Covalent Compounds
Covalent Bonding and the Periodic Table
Lewis structures are electron-dot structures for
molecules. They show the location of all valence e−.
113
Lewis Structures
A molecular formula shows the number and identity
of all of the atoms in a compound, but not which
atoms are bonded to each other.
A Lewis structure shows the connectivity between
atoms, as well as the location of all bonding and
nonbonding valence electrons.
General rules for drawing Lewis structures:
• Draw only valence electrons.
• Give every main group element (except H) an
octet of e−.
• Give each hydrogen two e−.
114
Lewis Structures
HOW TO Draw a Lewis Structure
Step [1]
Arrange the atoms next to each other that
you think are bonded together.
• Place H and halogens on the periphery, since they
can only form one bond.
H
For CH3Cl:
H C
Cl
H
115
Lewis Structures
HOW TO Draw a Lewis Structure
Step [2] Count the valence electrons.
• For main group elements, the number of valence
e− is equal to the group number.
• The sum gives the total number of e− that must
be used in the Lewis structure.
For CH3Cl:
1 C x 4 e− = 4 e−
3 H x 1 e− = 3 e−
1 Cl x 7 e− = 7 e−
14 total valence e−
116
Lewis Structures
HOW TO Draw a Lewis Structure
Step [3]
Arrange the electrons around the atoms.
• Place one bond (two e−) between every two atoms.
• Use all remaining electrons to fill octets with lone
pairs, beginning with atoms on the periphery.
117
Lewis Structures
HOW TO Draw a Lewis Structure
For CH3Cl:
H
H C
e−
2 on
each H
H
4 bonds x 2 e− = 8 e−
Cl
e−
8
on Cl
+ 3 lone pairs x 2 e− = 6 e−
14 e−
All valence e− have
been used.
118
Lewis Structures
Multiple Bonds
• When it is not possible to give every main group
element an octet by placing only single bonds in a
molecule, Lewis structures must contain one or
more multiple bonds.
• A double bond contains four electrons in two
two-electron bonds.
O
C
O
• A triple bond contains six electrons in three
two-electron bonds.
N
N
119
Molecular Shape
Valence Shell Electron Pair Repulsion
(VSEPR) Theory
• To determine the shape around a given atom,
first determine how many groups surround the
atom.
• A group is either an atom or a lone pair of
electrons.
• Use the VSEPR theory to determine the shape.
• The most stable arrangement keeps the groups
as far away from each other as possible.
120
Molecular Shape
Two Groups Around an Atom
• Any atom surrounded by only two groups is
linear and has a bond angle of 180o.
• An example is CO2:
• Ignore multiple bonds in predicting geometry.
Count only atoms and lone pairs.
121
Molecular Shape
Three Groups Around an Atom
• Any atom surrounded by three groups is
trigonal planar and has bond angles of 120o.
• An example is H2CO:
122
Molecular Shape
Four Groups Around an Atom
• Any atom surrounded by four groups is
tetrahedral and has bond angles of 109.5o.
• An example is CH4:
123
Molecular Shape
Four Groups Around an Atom
• If the four groups around the atom include one
lone pair, the geometry is a trigonal pyramid
with bond angles of ~109.5o.
• An example is NH3:
124
Molecular Shape
Four Groups Around an Atom
• If the four groups around the atom include two
lone pairs, the geometry is bent and the bond
angle is 105o (i.e., close to 109.5o).
• An example is H2O:
125
Molecular Shape
126
Electronegativity and Bond Polarity
• Electronegativity is a measure of an atom’s
attraction for e− in a bond.
• It tells how much a particular atom “wants” e−.
127
Electronegativity and Bond Polarity
• If the electronegativities of two bonded atoms
are equal or similar, the bond is nonpolar.
• The electrons in the bond are being shared
equally between the two atoms.
128
Electronegativity and Bond Polarity
• Bonding between atoms with different electronegativities yields a polar covalent bond or dipole.
• The electrons in the bond are unequally shared
between the C and the O.
• e− are pulled toward O, the more electronegative
element; this is indicated by the symbol δ−.
• e− are pulled away from C, the less electronegative
element; this is indicated by the symbol δ+.
129
Electronegativity and Bond Polarity
130
Polarity of Molecules
The classification of a molecule as polar or nonpolar
depends on:
• The polarity of the individual bonds.
• The overall shape of the molecule.
Nonpolar molecules generally have:
• No polar bonds.
• Individual bond dipoles that cancel.
Polar molecules generally have:
• Only one polar bond.
• Individual bond dipoles that do not cancel.
131
Polarity of Molecules
To determine the polarity of a molecule that has two
or more polar bonds:
1. Identify all polar bonds based on electronegativity.
2. Determine the shape around individual atoms.
3. Decide if individual dipoles cancel or reinforce.
132