Chapter 6
Bonding
Objectives
 Write the electron configuration of ions of representative
elements.
 Define ionic bonds in terms of the difference in
electronegativity of the atoms.
 Write the Lewis electron dot structures of atoms, ions and
ionic compounds.
 Write the name of an ionic compound (binary, polyatomic and
transition metal) given its formula, or its formula given its
name.
 Describe the properties of ionic, metallic and covalent
molecules and identify the forces holding them together.
Valence Electrons
• Mendeleev ordered his periodic table
according to chemical behavior.
• A pattern he saw was that elements in the
same group behaved the same chemically.
• This is due to its valence electrons.
• Valence electrons are the outermost
electrons and are involved in bonding.
Calculating valence electrons
• Groups on the periodic table can help!
What’s left out
on the periodic
table?
Transition Metals
• Transition metals can have different amounts of valence
electrons depending on how they bond with nonmetals
Decide how many valence
electrons each one has?
1.
2.
3.
4.
5.
6.
7.
Li
F
Kr
Mg
Group 13
Carbon
Phosphorus
1.
2.
3.
4.
5.
6.
7.
1 valence electron
7 valence electron
8 valence electron
2 valence electron
3 valence electron
4 valence electron
5 valence electron
Gilbert Lewis
• Lewis developed the
idea of the octet and
coined the term Lewis
dot structures
• He was nominated 35
times for the Nobel
Prize in chemistry, but
never won.
Electron Dot structures
• A diagram that shows only the valence electrons around
the atom as dots
Some rules:
• Each side can only have 2 dots, for a maximum amount
of 8 dots
• An Exception to group 18 is Helium, it will only have 2
Electron Dot structures
• Draw the correct Lewis dot structure for each given:
1 V.E
6 V.E
2 V.E
7 V.E
3 V.E
8 V.E
4 V.E
5 V.E
Try these on your own:
Electron dot structures
Think back to Nobel Gases
• Nobel gases are unreactive…
• What do you notice about their electron
dot structures??
Octet Rule
• Atoms want to achieve stability or noble gas
configuration by attaining 8 valence
electrons.
• Maximum number in the s and p orbitals.
• Atoms do not want to gain or lose more than
3 electrons when bonding!
Which is the only group that
has full octets?
• NOBLE GASES!
Ions
Ions
• Ions are formed when an atom loses or
gains electrons
• Ions can gain or lose just one or multiple
electrons to bond and achieve stability
• Ions do not want to gain or lose more than 3
electrons though when bonding
Cations
• Cations are formed when an atom loses
electrons and becomes positive
• Metals form cations
• Examples: magnesium, potassium,
aluminum
Anions
• Anions are formed when an atom gains
electrons and becomes negative
• Nonmetals form anions
• Examples: fluorine, oxygen, phosphorus
How do we know if atoms
lose or gain electrons?
• Let’s look at their Valence electrons
• Every atom wants to be at 8 valence electrons
• Sodium has 1 VE, so its easier to lose 1 and have a +1
• Chlorine has 7 VE, so its easier to gain 1 and would have
a–1
Charge or Oxidation state
• The charge corresponds with the amount of valence
electrons:
Did it lose or gain?
• Oxide ion
1. O-2: gained 2 electrons
• Cesium ion
2. Cs+1: lost 1
• Aluminum ion
3. Al+3 lost 3 electrons
• Bromide ion
4. Br-1 gained 1 electrons
• Phosphide ion
5. P-3 gained 3 electrons
• Valence electrons: Gain or lose, and how many?
1.
2.
3.
4.
5.
Strontium
Hydrogen
Sulfur
Xenon
Iodine
1.
2.
3.
4.
5.
Lose 2
Lose 1
Gain 2
Neither, neutral
Gain 1
Polyatomic Ions
Polyatomic Ions
Ions
Review
Predicting Ionic Charges
Group 1: Lose 1 electron to form 1+ ions
H+
Li+ Na+
K+
Predicting Ionic Charges
Group 2: Loses 2 electrons to form 2+ ions
Be2+
Mg2+
Ca2+
Sr2+
Ba2+
Predicting Ionic Charges
B3+
Al3+
Ga3+
Loses 3
electrons to form
3+ ions
Group 13:
Predicting Ionic Charges
N3- Nitride
P3- Phosphide
As3- Arsenide
Group 15:
Gains 3
electrons to form
3- ions
Predicting Ionic Charges
O2- Oxide
S2- Sulfide
Se2- Selenide
Gains 2
electrons to form
2- ions
Group 16:
Predicting Ionic Charges
F1- Fluoride
Br1- Bromide
Cl1-Chloride
I1- Iodide
Gains 1
electron to form
1- ions
Group 17:
Predicting Ionic Charges
Group 18:
Stable Noble gases
do not form ions!
Predicting Ionic Charges
Groups 3 - 12:
Many transition elements
have more than one possible oxidation state.
Iron(II) = Fe2+
Iron(III) = Fe3+
Predicting Ionic Charges
Groups 3 - 12:
Some transition elements
have only one possible oxidation state.
Zinc = Zn2+
Silver = Ag+
Chapter 6
Types of Bonding
• Atoms seldom exist in nature as
independent particles.
• Nearly all substances are made up of a
combination of atoms that are held
together by chemical bonds.
• Chemical Bond – a mutual electrical
attraction between the nuclei and
valence electrons of different atoms that
bind the atoms together.
When atoms bond, their valence electrons
are redistributed in ways that make the
atoms more stable.
The ways in which the electrons are
redistributed determines the type of bonding.
Metals tend to lose electrons to form positive
ions, or cations, and nonmetals tend to gain
electrons to form negative ions, or anions.
Chemical bonding that results from the
electrical attraction between cations and
anions is called ionic bonding.
Chemical bonding that results from the
sharing of electrons between atoms is
covalent bonding.
In covalent bonds the shared electrons are
owned equally by the two bonded atoms.
You can estimate the type of bonding (ionic or
covalent) between elements by the difference
in electronegativities (Chapter 5).
Electronegativity
• This chart will help you determine if it is polar or nonpolar
0.0-0.3 Non polar covalent
0.3-1.7 polar covalent
>/= 1.8 Ionic
Polar Covalent vs. Non-polar Covalent
Blue shading represents electron density
Sample Problem
Page 163
Use electronegativity differences and Figure
6-2 to classify bonding between sulfur, S, and
the following elements: hydrogen, H; cesium,
Cs; and chlorine, Cl as either ionic, polar
covalent or non-polar covalent.
electronegativity
difference
bond type
S and H
2.5 - 2.1 = 0.4
polar
S and Cs
2.5 - 0.7 = 1.8
ionic
S and Cl
3.0 - 2.5 = 0.5
polar
Ionic Bonds
Opposites attract!
• An ionic compound is composed of positive
and negative ions that are combined so that
the charges are equal.
• Cations will combine with anions to form
ionic compounds or salts. (metal and
nonmetal)
• This electrostatic force that holds them
together is called an ionic bond!
Example: sodium chloride consists of a
positive ion (Na) with a +1 charge and
negative ion (Cl) with a -1 charge.
Na+ + Cl-
NaCl
Draw electron configurations and
Lewis Dot Structures.
Ionic Bonds
• Cation + Anion  Ionic bond (Ionic compound)
• Salt (NaCl) Ionic bond between a metal and non-metal
• Look at the electron dot structures.
• Note how the sodium donates its electron to chlorine,
now chlorine has an octet and a negative charge
Structure of Sodium Chloride Crystal
Structure of Sodium Chloride Crystal
Ionic bonds
• Show the ionic bonds of the following using the electron
dot structure method.
1. Magnesium and oxygen
(MgO)
2. Potassium and Sulfur
(K2S)
3. Calcium and chloride
(CaCl2)
Properties of
ionic compounds
Properties of ionic compounds
• Electrically neutral
• Hard, but brittle
• Most form crystal
lattices at room
temperature
• Generally, have high
melting points and
Boiling Points
Fluorite
(CaF2)
Cinnabar
(HgS)
Properties of ionic compounds
• Conduct electricity
when dissolved in
water or as a liquid.
• Solids do not conduct
electricity.
Ionic Compounds in Water
• When ionic
compounds are
placed in water,
they will dissociate.
• This means the
anions and cations
will split apart and
form weak bonds
with the water
molecules.
Sodium chloride vs sugar
NaCl
• Sodium is a metal
• Chloride is a nonmetal
• Melting point ~800oC
C12H22O11
• C, H, and O are
nonmetals
• Melting point ~185oC
Which one (or both) is/are an ionic
compound(s)?
Crystal Lattice
• Ionic compounds form in
repeating patterns of
anions and cations.
• Their crystal structure
will be the same for a
particular compound.
Ionic crystals
• Ionic compounds that are crystals are made out of small
pieces called unit cells
• A unit cell is the simplest repeating unit in a crystal
• NaCl’s unit cell:
Lesson check:
True or False:
1. Ionic compounds
have low melting
points
2. Nitrogen and oxygen
form an ionic
compound
3. Ionic Compounds are
normally liquids
4. Ionic Compounds
conduct electricity
when dissolved
1.
2.
3.
4.
False
False
False
True
Metallic Bonding
• What are some properties of
metals?
• Think back to Chapter 1
Properties of metals
• Good conductors of
electricity
• Ability to be drawn
into a wire (Ductile)
• Ability to be
hammered, without
breaking (Malleable)
• Not brittle
Metallic bonds
• Metals are composed of closely packed cations held
together by their outer electrons.
• These outer electrons can be referred to as a sea of
electrons. The electrons move freely between metal
atoms.
• This “sea of electrons” is what allows metals to be
molded, hammered, or stretched.
Metallic bonding
• Metallic Bonds are the forces or attraction between
those free floating outside electrons and the positively
charged metal ions
Metal alloys
• An alloy is a mixture of 2 or more elements
(one must be a metal)
• These are uniform throughout, so a
homogeneous mixture
• Examples: Brass (copper and zinc); Sterling
silver (silver and copper); Bronze (copper
and tin)
Why have alloys?
• Alloys are important because they are combining
properties and are often superior compared to the pure
elements
• Typically, more inexpensive than the pure element:
• Sterling silver vs pure silver
$0.95 vs $1.68
Think about this:
• A bronze statue is beginning to turn green; bronze is an
alloy made of copper and tin
• Which element is causing it to become green?
•
Hint: Think statue of liberty
Covalent
Molecules
Molecules and Molecular
Compounds
• Ionic bonds are a + and -, but what about CO2?
• What is type of elements are C and O?
• When nonmetals bond together a covalent bond is
created and we call them molecules or molecular
compounds!
Molecules
• Molecules are neutral atoms that are joined together by
covalent bonds
• Molecular Compound another way a saying molecule
• Molecular formula shows you how many atoms of each
element is in a substance
• Example: CO2 , NH4
Octet Rule and Covalent
Bonding
• An octet is 8 valence electrons and want to achieve
noble gas configuration!
• Molecules want the same thing, but they share their
valence electrons
Sharing electrons
• Recall that ionic bonds give and take electrons…
• Molecules share their electrons between the 2 atoms.
• When they share their valence electrons, a covalent
bond is made
Single covalent bonds
• When atoms share one pair of electrons they form a
single covalent bond
• Example: H2
• Let’s draw it:
Show these diatomics:
• Cl2
• Br2
• I2
• F2
• What about H2O?
Structural Formula
• Electron dot
structure
represents bonds
as 2 dots coming
together:
• A structural formula
represents covalent
bonds as dashes
What did we call those 2 dots
next to one another?
• Lone pairs or
unshared pairs
• They do NOT
participate in
bonding, but you
must show them!
Try these on your own:
• NH3
• CH4
• H2O2
• PCl3
Double Bonds
• Atoms that share two pairs of electrons
• Example: CO2
Double Bonds
• Draw O2:
Triple Bonds
• Atoms that share three pairs of electrons:
• Example: N2
Properties of
Covalent
Molecules
Properties of covalent molecules
• Made out of
nonmetals
• Can be a solid, liquid,
or gas at room
temperature
• Low melting point and
boiling points
• Poor to nonconductors of heat
and electricity
H2O vs NaCl
Liquid water
Solid water
Strengths of
covalent bonds vs.
ionic bonds
Bonding Theories
How do we decide where to
find electrons?
• The modern atomic theory tells us that they are most
probable at certain locations
Molecular Orbitals
• When two atoms combine, their atomic orbitals combine
and overlap to produce molecular orbitals
• A molecular orbital belongs to the whole of the molecule
• Each orbital can only contain 2 electrons
• When an orbital overlaps and participates in a covalent
bond, it can be classified as a bonding orbital
Sigma bonding (σ)
• The first bond between sharing atoms is classified as
sigma bonds
Pi bonding (π)
• The second type of bonding can be a pi
• Remember the first is a sigma and the rest can be pi
bonds
Molecular Geometry
VSEPR Theory
VSEPR Theory
• Valence-Shell-Electron-Pair-Repulsion theory
• This theory helps us understand the 3D structure of
molecules and their properties.
• Bonding and unshared pairs of valence electrons
become very important to us within VSEPR theory!
• The shapes of molecules are determined because
electron pairs want to be far apart from each other
(repulsion).
AXE
– Method to represent compounds
A represents the central atom
X represents the bonding atoms
E represents the lone pairs the central atom has
Compounds with
no lone pairs
• Draw or build CO2
• Meaning 1 central atom, 2 bonded atoms
• It has a linear shape
• No lone pairs
• A bond angle of 180o
• Bonding pairs are far apart from each other
• Draw or build BF3
• This has a trigonal planar
• Meaning 1 central, 3 bonded
• Bond angle: 120o
• Bonds pointing to the corners of a triangle
• Draw or build CH4
• This has a tetrahedral
• AX4
• Meaning 1 central, 4 bonded
• Bond angle: 109.5o
Problem
Use VSEPR theory to predict the
shape of:
aluminum trichloride, AlCl3
hydrogen iodide, HI
carbon tetrabromide, CBr4
dichloromethane, CH2Cl2
Compounds with
lone pairs
Lone pairs occupy space, but
only bonded atoms determine
the name
• Draw or build H2O
• This has a bent shape
• AX2E2
• Meaning 1 central, 2 bonded, 2 lone pairs
• Bond angle: 105o
• Similar angles to the
tetrahedral bond angles
• Draw or build NH3
• This has a trigonal pyramidal
• AX3E
• Meaning 1 central, 3 bonded, 1 lone pair
• Bond angle: 107o
• Similar angles to the
tetrahedral bond angles
What shape is each one?
• BeCl2
• Linear
• OF2
• Bent
• AlCl3
• Trigonal Planar
• PCl3
• Trigonal pyramidal
• CF4
• Tetrahedral
Bond Polarity
Switch presentations – slide 80
Bond polarity
• Since, atoms are sharing within a covalent bond…
• If they share equally they are a nonpolar covalent bond
• Examples: Diatomic atoms are nonpolar because they
pull on each other evenly
Bond polarity
• Since, atoms are sharing within a covalent bond…
• If they share unequally they are a polar covalent bond
• Examples: HCl, H2O
Bond polarity
• Since, polar bonds are unequally sharing we will have
dipoles.
• But how will we decide polarity??
• Electronegativity!!
Dipoles
• The more electronegative will have the arrow point
towards it and have a slightly negative charge
• The less electronegative will have a slightly positive
charge
Electronegativity
• This chart will help you determine if it is polar or nonpolar
0.0-0.4 Non polar covalent
0.4-1.0 Slightly polar covalent
1.0-2.0 Very polar covalent
>/= 2.0 Ionic
Decide the polarity of the
following:
• N-H
• 0.9 slightly polar
• F-F
• 0 Nonpolar
• Ca- Cl
• 2.0 ionic
• Al- Cl
• 1.5 very polar