UNIT 2: Structure and Properties of Matter
Chapter 3: Atomic Models and
Properties of Atoms
Chapter 4: Chemical Bonding and
Properties of Matter
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter
Chapter 4: Chemical Bonding
and Properties of Matter
The chemical bonding in a
substance influences the
shape of its molecules, and
molecular shape influences
the properties of that
substance. One of the
properties of iron is its
strength, which makes it
ideal for use in support
structures.
The strength of iron makes it useful in
items such as horseshoes.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
4.1 Models of Chemical Bonding
Three types of chemical bonding are ionic, covalent,
and metallic.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Electronegativity
Electronegativity is the relative ability of an atom to
attract shared electrons in a chemical bond.
TO PREVIOUS
SLIDE
What general trends in electronegativity are
shown in the periodic table?
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Electron Sharing and Electronegativity
Electronegativity difference, ΔEN, between two atoms bonded
together can be low, intermediate, or high. The electron density
diagrams below show the differences in the bonds.
• when ΔEN is 0: electrons are equally shared
• when ΔEN is 1: electrons are more closely associated with
the more electronegative atom
• when ΔEN is high, there is little sharing of electrons
TO PREVIOUS
SLIDE
Bonding is a continuum between equal sharing and
minimal sharing of electrons.
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Electron Sharing and Electronegativity
Scientists have categorized types of bonds according to ΔEN.
• ΔEN between 1.7 and 3.3:
mostly ionic
• ΔEN between 0.4 and 1.7:
polar covalent
• ΔEN between 0.0 and 0.4:
mostly covalent (non-polar)
TO PREVIOUS
SLIDE
Three categories of bonds have
been set based on ΔEN .
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Metallic Bonding
Chemists use the electron-sea model to describe metallic
bonding. The model proposes that the valence electrons of
metal atoms move freely among the ions, forming a “sea” of
delocalized electrons that hold the metal ions rigidly in place.
Microscopic analysis shows that the
structure of metals consists of aggregates
of crystals.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Properties of Metals
Melting and Boiling Points
• the stronger the bonding forces, the higher the melting and
boiling points of pure metals
TO PREVIOUS
SLIDE
Periodic table trends include:
1. For Group 1, melting points decrease as the atomic number increases.
2. For Groups 1 to 6, across a period, melting points increase as atomic number increases.
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Properties of Metals
Electrical and Thermal Conductivity
• Metals are good conductors because their electrons are
free to move from one atom to the next.
Malleability and Ductility
• Based on the electron-sea model, metals can be shaped
because, when struck, the metal ions can slide by one
another while the electrons still surround them.
Hardness
• The variation between metals is due to differences in
crystal size (smaller ones make harder metals).
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Alloys
Alloys are solid mixtures of two or more metals.
• the addition of the second metal, even in a very small amount,
can significantly affect the properties of a substance
• in some cases, non-metal atoms, such as carbon, are added
TO PREVIOUS
SLIDE
If atoms of the second
metal are similar in size to
the first metal, they take the
place of those atoms.
If atoms of the second metal are
much smaller than atoms of the first
metal, they will fit in spaces
between the larger atoms.
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Ionic Bonding
• occurs when ΔEN is between 1.7 and 3.3
• essentially, involves one atom losing one or more
electrons and another atom gaining those electron(s)
There are different ways to show the transfer of electrons in the
formation of ionic compounds.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Ionic Crystals
Ionic compounds exist as crystal lattice structures with
particular patterns of alternating positive and negative ions.
The unit cell is the smallest group of ions that is repeated.
NaCl forms a cubic
crystal lattice structure.
Different types of crystal structures can form.
• the relative sizes and charges of the ions affect the type of
crystal structure that an ionic compound will form.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Properties of Ionic Compounds
Melting and Boiling Points
• high due to very strong attractions between ions
Solubility
• ionic compounds are soluble in water when the attractive
forces between the ions and water molecules are stronger
than the attractive forces among the ions themselves
TO PREVIOUS
SLIDE
When sodium chloride (NaCl)
dissolves in water, attractive
forces between water
molecules and NaCl ions act
to break apart the ionic bonds.
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Properties of Ionic Compounds
Mechanical Properties
• hard and brittle, so will break apart when struck
Ionic crystal will break on
smooth planes, where like
charges become aligned.
Conductivity
• solids do not conduct because ions cannot move
• compounds conduct when dissolved in water and ions can
move
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Covalent Bonding
• occurs when ΔEN is less than 1.7
• covalent bonds are classified into two types:
• polar covalent: atoms do not share electrons equally
• non-polar covalent: atoms share electrons almost equally
Forces in covalent bonds:
• both attractive and repulsive forces play a role
The length of a covalent bond
is determined by different
electrostatic forces.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
LEARNING CHECK
Describe the chemical bonding and
structure of NaCl. How do bonding and
structure influence the general properties
of the substance?
TO PREVIOUS
SLIDE
Answer on
the next slide
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
LEARNING CHECK
NaCl is composed of a metal atom bonded to
a non-metal atom with ΔEN > 1.7. As such, the
bond is classified as ionic. It exists as a cubic
crystal lattice structure, with an alternating
pattern of chloride ions and sodium ions.
Properties of NaCl include high melting and
boiling points; solubility in water; hard and
brittle; a poor conductor as a solid, but it does
conduct electricity when dissolved in water.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Quantum Mechanics and Bonding
Quantum mechanics is used to explain and describe chemical
bonding. It is also used to account for shapes of molecules.
Valence Bond (VB) Theory explains bond formation and
molecular shapes based on orbital overlap.
• The region of overlap has a maximum capacity of two
electrons, which have opposite spins.
• There should be maximum overlap of orbitals, since the
greater the overlap, the stronger and more stable the bond.
• Atomic orbital hybridization is used to help explain the
shapes of some molecules.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Quantum Mechanics and Bonding
Molecular Orbital (MO) Theory explains bond
formation and molecular shapes based on the formation of
new molecular orbitals.
According to MO theory:
• Covalent bond formation involves atomic orbital overlap
that results in formation of new orbitals called molecular
orbitals.
• Molecular orbitals have shapes and energy levels that are
different from those of atomic orbitals.
• The electrons in molecular orbitals are delocalized
throughout the orbital.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Explaining Single Bonds
For molecules like hydrogen fluoride:
• the 1s orbital of H overlaps with the half-filled 2p orbital
of F
According to MO theory, the bond is a sigma (σ) bond,
which is symmetrical and freely rotates.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Explaining Single Bonds
For molecules like methane:
• the VB theory of hybrid orbitals is used to explain
molecular shape
• carbon forms four hybrid orbitals (sp3) by combining
three 2p orbitals and a 2s orbital so that four identical
bonds can be created
The four sp3 orbitals of C overlap with the s orbitals of H to form methane.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Explaining Double Bonds
Hybrid orbitals are used to
explain the structure of ethene
or molecules like ethene.
TO PREVIOUS
SLIDE
• it is planar with ~120º bond angles
• the structure is explained by formation of 3 sp2 hybrid orbitals
for each carbon (a 2s orbital mixes with two 2p orbitals)
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Explaining Double Bonds
For bond formation in ethene:
• one sp2 orbital of each carbon overlaps to form a σ bond
between the carbons
• two sp2 orbitals of each carbon overlap with the 1s
orbitals of the hydrogens to form σ bonds
• the lobes of the 2p orbitals of each carbon overlap above
and below the plane to form a pi (π) bond
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Explaining Triple Bonds
For molecules like ethyne:
• the linear structure is explained by formation of 2 sp
hybrid orbitals for each carbon (a 2s orbital + a 2p
orbital)
• sigma bonds form from overlap between sp of each
carbon and between sp of carbons and 1s of hydrogens
• two pi bonds form from overlap of the two 2p orbitals of
each carbon
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Types of Hybridization
• The names of hybrid orbitals (formed by the combination of two
or more orbitals in the valance shell of an atom) indicate the
number and types of atomic orbitals that were combined.
• Atoms of Period 3 elements can have d orbital hybridization with s
and p orbitals.
• The number of hybrid orbitals that form is the same as the number
of atomic orbitals that are combined.
TO PREVIOUS
SLIDE
Each hybrid orbital has a certain overall shape.
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Allotropes
Allotropes are compounds that consist of the same element
but have different physical properties.
An example is allotropes of carbon, which differ in the
pattern of covalent bonds between carbon atoms.
TO PREVIOUS
SLIDE
Allotropes of carbon: A graphite, B diamond,
C buckyballs, D nanotubes
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Covalent Network Solids
Network solids are substances that consist of atoms bonded
covalently in a continuous two- or three-dimensional array.
There is no natural beginning or end to the chains of atoms.
Silicon dioxide, SiO2, exists as a
network solid that is represented as
(SiO2)n.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.1
Section 4.1 Review
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
4.2 Shapes, Intermolecular Forces, and
Properties of Molecules
Molecular compounds form a much greater variety of
structures than ionic compounds form.
Understanding the properties of molecules requires an
understanding of their three-dimensional shapes.
Different theories and models are used to predict molecular
shapes.
The shape of a molecule is the result
of the presence of atoms, bonding
electrons, and non-bonding electrons,
as well as forces of attraction and
repulsion.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Depicting Two-Dimensional Structures of
Molecules with Lewis Structures
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Some Exceptions When Drawing
Lewis Structures
Co-ordinate Covalent Bonds:
• one atom contributes both electrons
• bonds behave the same way as other
covalent bonds and therefore are not
indicated in Lewis structures
Expanded Octet (Expanded Valence):
• central atom has more than an octet of
electrons
• a feature of some Period 3 and higher
elements
TO PREVIOUS
SLIDE
The ammonium ion has a
co-ordinate covalent bond.
For SF6(g), 12 electrons are
around the central atom.
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Some Exceptions When Drawing
Lewis Structures
An Incomplete Octet:
• central atom has fewer than an octet
of electrons
Resonance Structures:
• measured bond lengths may not
support Lewis structures
• one of two or more Lewis structures
that show same relative position of
atoms but different positions of
electron pairs
TO PREVIOUS
SLIDE
In BF3(g), boron has an
incomplete octet.
Actual bond lengths in ozone
are between those of single
and double bonds.
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Predicting the Shapes of Molecules
Using VSEPR Theory
The valence-shell electron pair repulsion (VSEPR) theory
• is a model used to predict molecular shape
• is based on electron groups around a central atom being
positioned as far apart as possible (repulsion)
• predicts certain arrangements of electron groups
TO PREVIOUS
SLIDE
For VSEPR, there are five electron-group arrangements.
(Electron groups are represented by bars).
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Electron Groups and Molecular Shapes
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Electron Groups and Molecular Shapes
If one or more electron groups around a central atom is a
lone pair, different strengths of repulsive forces will alter
bond angles to differing degrees.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Summarizing Molecular Shapes
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Summarizing Molecular Shapes
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Guidelines for Using VSEPR Theory to
Predict Molecular Shape
TO PREVIOUS
SLIDE
UNIT 2 Chapter 3: Atomic Models and Properties of Atoms
Section 3.2
LEARNING CHECK
What is the electron-group
arrangement and molecular
shape of HCN?
TO PREVIOUS
SLIDE
Answer on
the next slide
UNIT 2 Chapter 3: Atomic Models and Properties of Atoms
LEARNING CHECK
HCN has two bonding groups and
no lone pairs.
The electron-group arrangement is
linear, and the shape of the
molecule is also linear.
TO PREVIOUS
SLIDE
Section 3.2
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Determining the Hybridization of
the Central Atom of a Molecule or Ion
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
LEARNING CHECK
What is the hybridization for
the phosphorus atom in the
molecule below?
TO PREVIOUS
SLIDE
Answer on
the next slide
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
LEARNING CHECK
The electron-group arrangement for
phosphorus is tetrahedral.
Therefore, P has an sp3 hybridization.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
The Influence of Molecular Shape
on Polarity
The shape of a molecule affects that molecule’s polarity.
• polar bonds have a bond dipole
• bond dipoles are indicated using vectors that point in the
direction of higher electron density
In a polar covalent bond, a partial positive charge is
associated with one atom and a partial negative
charge is associated with the other atom.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Determining Whether a
Molecule is Polar
A molecule with one or more polar bonds is not necessarily a
polar molecule. The molecule’s shape must be considered. The
polarity as a whole can be determined by adding the vectors.
TO PREVIOUS
SLIDE
Both water and carbon dioxide have two polar bonds. But water’s bent shape results in a
polar molecule, while carbon dioxide’s linear shape results in a non-polar molecule.
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Molecular Shapes and Polarities
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
How Intermolecular Forces Affect the
Properties of Solids and Liquids
Intermolecular forces exist between ions and molecules
and influence the physical properties of substances.
Categories of forces:
• dipole-dipole
• ion-dipole
TO PREVIOUS
SLIDE
• induced dipole
• dispersion
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Dipole-Dipole
Dipole-dipole forces:
• are forces of attraction between polar molecules, which
have a region of partial positive charge and a region of
partial negative charge
• are a main reason for melting and boiling point
differences between polar and non-polar molecules
• include hydrogen bonding, as an example of one type
Hydrogen bonding
(dotted lines) in water
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Ion-Dipole
Ion-dipole forces:
• are forces of attraction between partial charges on polar
molecules and ions
• depend on the size and charge of the ion and the
magnitude of the partial charge and size of the molecule
• are involved in the process of hydration
Ion-dipole intermolecular forces.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Induced Dipoles
Dipole-induced dipole forces:
• are forces of attraction between a polar molecule and a
non-polar molecule that has an induced (temporary)
dipole due to the nearby polar molecule
Ion-induced dipole forces:
• are forces of attraction between an ion and a non-polar
molecule that has an induced dipole due to the nearby
ion
A dipole can be induced
in a non-polar molecule.
TO PREVIOUS
SLIDE
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Dispersion Forces
Dispersion forces:
• are forces of attraction between all molecules, including
non-polar molecules
• are due to spontaneous temporary dipoles that form due to
the constant motion of electrons in covalent bonds
• depend on the size and shape of the molecules
• the larger and more linear the molecule, the greater the
force of attraction
TO PREVIOUS
SLIDE
The more linear molecule has a
higher boiling point because the
dispersion forces are greater.
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Section 4.2 Review
TO PREVIOUS
SLIDE