BONDING
IB TOPIC 4
Chapter 2
REVIEW OF ELEMENTS AND
COMPOUNDS
WHEN YOU SEE
EXERCISES
COPY OR DO THE
MEANS IMPORTANT TO KNOW even if you
do not copy.
WHEN YOU SEE
CHECK YOUR
HANDOUTS
ELEMENTS




Elements are pure
substances that can not
be further broken down
chemically or physically.
Elements only contain one
type of atom.
There are about 116
different kinds of
elements; 92 are
considered to be natural
and the remaining are
synthesized.
All the elements are listed
on the periodic table.
METALS VS. NON-METALS
Metals (Na, Ca Al)
Elements
Non-metals (Cl, Si,
Ar)
Substances are grouped (classified) according to their
physical and chemical properties.
CHARACTERISTICS OF METALS
Physical Properties





Solids at standard conditions
(SATP) except Hg which is a
liquid
Malleable
Ductile
Good conductors of heat and
electricity
Grey and shiny; reflect light
when polished
Chemical Properties





Are very reactive
React vigorously with water
Form ionic compounds with
non-metals
Combine with chlorine to from
a white solid that dissolves in
water
Combine with O2 to form metal
oxide
TERMS
malleable - can be hammered into shapes or bent
without breaking
 ductile - can be drawn into wires

NON-METALS
Physical Properties
Chemical Properties
Brittle and dull if a solid
 Poor conductors of heat
and electricity
 Can be solids, liquids or
gases at SATP
 Gases: the noble
gases, H2, O2, N2, Cl2,
F2
 Liquids: Br2
 solids: remainder


Can be very reactive (F2 is
most reactive)
 Do not react with acids
 Form ionic compounds
with metals
 Form molecular
compounds (covalent
bonds) with other nonmetals
COMPOUNDS
 Compounds
are pure substances that contain two
or more types of atoms that combine (BOND) to
chemically to form a new substance that has its
own unique chemical and physical properties.
 A chemical equation representing two elements in
a chemical reaction that results in the formation of
a new substance, a compound:
4Fe
iron
+ 3O2
+ oxygen
→
element + element
→
2Fe2O3
iron oxide (rust)
compound
COMPOUNDS
 They
can be broken down into a simpler
type of matter (elements) by chemical
means (but not by physical means).




Atoms neutral. They are constantly moving. When two atoms
approach two sets of forces are present:
1. Repulsive forces between the electrons of the two atoms
and between the protons of the atoms
2. Attractive forces (electrostatic attraction) between the
protons (+) of one atom and the valence electrons (-) of
another.
When the attractive forces are greater than the repulsive
forces, a chemical bond occurs.
Atoms wish to achieve the stability of noble gases by filling
their valence shells by exchanging or sharing electrons- i.e.
forming bonds
Bonds do not form unless it results in greater stability
(usually).

Bond dissociation energy (page 27) is the energy
required to break the bonds between 1 mol (see later for
definition of this unit) of bonded atoms. The larger the
value, the stronger the bond.

It takes energy to break bonds
(energy in = endothermic)

But when bonds are formed, energy is released
(energy out = exothermic).
3 Types of
Chemical Bonds
Ionic
METAL + NONMETAL
(section 2.1)
Covalent
NONMETAL +
NONMETAL
(section 2.3)
Metallic
METAL + METAL
(section 2.2)
IONIC BONDING
Page 28
4.1 IONIC BONDING
4.1.1
4.1.2
4.1.3
4.1.4
4.1.5
4.1.6
4.1.7
4.1.8
Describe the ionic bond as the electrostatic attraction between
oppositely charged ions.
Describe how ions can be formed as a result of electron transfer.
Deduce which will be formed when elements in groups 1, 2 and
3 lose electrons.
Deduce which ions will be formed when elements in groups 5, 6,
and 7 gain electrons.
State that transition elements can from more than one ion.
Include examples such as Fe2+ and Fe3+
Predict whether a compound of two elements would be ionic
from the position of the elements in the periodic table or from
their electronegativity values.
State the formula of common polyatomic ions formed by nonmetals in periods 2 and 3
Examples include NO3¯, OH¯, SO42¯, CO32¯, PO43¯, HCO3¯, NH4+
Describe the lattice structure of ionic compounds
Be able to describe the structure of NaCl as an example
4.1.1-2



Formation of NaCl
An ionic bond is an electrostatic attraction between two
oppositely charged ions.
Forms between metals and non-metals, or metals and
polyatomic ions. The transition metals may have multiple
charges.
Electrons are exchanged in ionic bonds: metals lose electrons
while non-metals gain electrons.
4.1.3-4
1
2
3
4
5
6
7
0
loses
1e-
loses
2e-
loses
3e-
(shares)
gains
3e-
gains
2e-
gains
1e-
NA
1+
2+
3+
NA
3‾
2‾
1‾
NA
cation
cation
cation
anion
anion
anion
Na+
Mg2+
Al3+
P3-
S2-
Cl-
C
NAMING IONS
Metallic ions use the elemental
name plus the word ion
Mg 2+
Al3+
magnesium ion
aluminium ion
Non-metallic ions end in ide
Cl ¯ chlorine = chloride
S2¯ sulfur = sulfide
Ar
OTHER IONS
In addition to simple monatomic ions (Mg2+, F- ,
O2-, Al3+) ions can be:
1.
Multiple charged (Fe3+, Fe 2+)
-all are cations
- use the Roman numeral in the name to indicate which ion
is used. e.g. iron (III) oxide = Fe2O3
2. Polyatomic (NO3¯, OH¯, SO42¯)
-all are anions except NH4+
-names ends in ite, ate, or ide
e.g. Mg(NO3)2 - magnesium nitrate
POLYATOMIC IONS
 polyatomic
ions are tightly bound groups of atoms
that behave as a unit


 All
there is one charge for the entire unit
Example: OH‾ hydroxide ion; O and H act together;
overall charge is 1‾
polyatomic ions will have a negative charge
EXCEPT for ammonia, NH4+
 Memorize these polyatomic ions and their charges:
NO3¯, OH¯, SO42¯, CO32¯, PO43¯, HCO3¯, NH4+
 Ionic
compounds are formed when a cation and
anion bond in such a manner to cancel their
charges- i.e. ionic compounds are neutral.
 Magnesium and chloride combine in a 1:2 ratio
and the formula is written as MgCl2
WRITING FORMULAS FOR IONIC COMPOUNDS
This formula unit is
for one unit of
magnesium chloride
Mg is the more
metallic element and
is written first.
MgCl2
Cl is the LESS
metallic and its
symbol is written
last
The subscript indicates that there is a ratio of two Cl
atoms to one Mg atom.
Use empirical formulas (lowest whole number ratio).
STEPS IN WRITING FORMULAS
 Refer
to page 7 of the Nomenclature Handbook.
This can be found on the blog.
WRITING FORMULAS FOR IONIC COMPOUNDS
CONTAINING POLYATOMIC IONS
NAMING WITH IUPAC RULES
ASSIGNED WORK

Complete page 5, 6 & 9 of Nomenclature Handbook
PROPERTIES
OF IONIC
COMPOUNDS
4.1.5




Ions that make up ionic compounds arrange themselves into a
regular 3-D structure called a lattice structure (or an ionic lattice or
crystal lattice).
The lattice`s geometry may vary depending on the size of the atoms
but it always involves a fixed arrangement of ions based on a
repeating unit.
The lattice can grow indefinitely and contains no fixed number of
ions, but the ratio is always the same for a given compound and this
is expressed in the empirical formula or formula unit.
DO NOT used the term molecule when describing ionic compounds.

Each Na+ is surrounded by 6 Cl- , and each Cl- is
surrounded by 6 Na+




Lattice energy or lattice enthalpy is the
energy required to separate 1 mol of an
ionic solid into its gaseous ions.
It indicates the strength of ionic
interactions and influences melting
point, hardness and solubility.
Ionic compounds have very high boiling
points because the ions must have
enough energy to break free from the
surroundings ions.
Ionic compounds are solid at room
temperature also because of their high
melting points.
Ionic compounds break rather than bend when they are struck.
The positive
and negative
ions in the
crystal are
arranged to
maximize
their
attractions.
When hit with enough force,
like charges move near each
other, and the repulsions
crack the piece apart.
Light
not on
In order to be
conductive, the
electrons or ions of a
substance must be
able to move freely.
Ionic compounds
conduct electricity
when in liquid form
or when dissolved in
water; BUT NOT
WHEN SOLID since
the ions are not free
to move in a solid.
•Ionic compounds have very high boiling points because the ions must
have enough energy to break free from the surroundings ions of the
liquid in order to become a gas.
•Because of this ionic compounds usually vaporize as ion pairs.
•Water is a polar substance (with positive and negative charged parts of the
molecule); hydrogen atoms are positively charged and the oxygen it is
negatively charged due to unequal distribution of the electrons (which are
negative).
•Water molecules bond to, and surround, the ions; the H-end is attracted to
the anion (-) while the O-end is attracted to the cations (+). This is called
dissolution (the act of dissolving).
All ionic compounds are soluble to some extent.
 However, the stronger the lattice enthalpy, the harder it
is for the water molecules to pull the salt ions away from
each other and the less soluble the substance.

KCl Period 4
NaCl Period 3
(Na is smaller
ion –higher
lattice energy
and less soluble)
Property
Ionic Compounds
Representative unit
Formula Unit
Bond formation
electrostatic attraction of oppositely
charged ions; anions accept electrons
that are donated by cations.
Type of elements that form this
bond
Metals + non-metals
State at room temp.
Crystalline solid (brittle)
Melting point
High (usually >3000C)
Electrical conductivity
Good conductor -as a liquid or
dissolved in water (but not as a solid)
Solubility in water
Most – highly soluble
Volatility
Not very volatile-heat needs to
overcome lattice energy
SAMPLE QUESTION
WHICH EQUATION REPRESENTS THE LATTICE ENTHALPY OF
MAGNESIUM OXIDE?
A.
B.
C.
D.
MgO(s) → Mg(s) + O2(g)
MgO(g) → Mg2+(g) + O2–(g)
MgO(s) →Mg2+(g) + O2(g)
MgO(s) →Mg2+(g) + O2-(g)
SAMPLE QUESTION
WHICH EQUATION REPRESENTS THE LATTICE ENTHALPY OF
MAGNESIUM OXIDE?
A.
B.
C.
D.
MgO(s) → Mg(s) + O2(g)
MgO(g) → Mg2+(g) + O2–(g)
MgO(s) →Mg2+(g) + O2(g)
MgO(s) →Mg2+(g) + O2-(g)
THE METALLIC BOND
Page 34
4.4
METALLIC BONDING
4.4.1 Describe the metallic bond as the electrostatic attraction
between a lattice of positive ions and delocalised electrons
4.4.2 Explain the electrical conductivity and malleability of metals



Metals conduct electricity and heat.
They are lustrous, malleable (able to be bent) and ductile
(able to be drawn in wires).
Metal atoms are not held together with covalent or ionic;
metals tend to donate electrons. But when a metal is
surrounded by other metal atoms, who will accept the
electron?
 Valence electrons on metals are held loosely and free to move


from one atom to the next (this explains conductivity of metals).
When the metal lets its electrons go, it is really forming a cation
which shares the electrons with all the other cations.
Metallic Bond – the electrostatic attractive force between a lattice
of cations surrounded by a sea of delocalized electrons (i.e. they
are confined to a particular location, but are free to move
throughout the structure).

When the metal is bent, the constantly moving sea of electrons
readjusts flowing in between the positive ions and preventing
them from coming into contact with each other. (If they did come
into contact with each other and repel, the substance would
fracture.)

When the metal is connected to a power supply, electrons enter
one end of the metal (from the negative end of the power supply)
and the same number of electrons leave the other end of the
metal (moving towards the positive end of the power supply).
Property
Ionic Compounds
Metals
Representative Unit
Formula Unit
NaCl CaF2
Element Symbol
Hg, Cu, Fe
Bond Formation
Electrostatic attraction of
oppositely charged ions
Protons in a sea
of delocalized electrons
Type of Elements
Metal + Non-metal
Metal + Metal
State at Room Temp.
Crystalline solid; hard and
brittle
Soft to hard solids; shiny,
malleable ductile.
Melting Point
High (usually >3000C)
Low to high
Electrical Conductivity Good conductor as liquid
or dissolved in water
Excellent conductor
as a solid or molten
Solubility in water
Not soluble
Most - high
PRACTICE
T or F Metals will conduct electricity when solid. T
 T or F Ionic compounds will conduct electricity when
solid. F
 What is responsible for the electrical conductivity of
metals? delocalized electrons
 Which is NaI, I2, and Na?

Na
NaI
What is the most reactive metal? Fr
 T or F Mg has a larger radius than Ba. F

I2
ASSIGNED WORK

Section 2.2 Exercises 1-6, page 36
COVALENT BOND
Page 36
4.2
COVALENT BONDING
4.2.1
Describe the covalent bond as the electrostatic attraction between a
pair of electrons and positively charged nuclei.
Single and multiple bonds should be considered. Examples should
include O2, N2, CO2, HCN, C2H4 (ethene) and C2H2 (ethyne).
4.2.2
Describe how the covalent bond is formed as the result of electron
sharing
4.2.3
Deduce the Lewis structure diagram (electron dot) structures of
molecules and ions for up to four electron pairs on each atom.
A pair of electrons can be represented by dots and crosses or a
combination, but not as Cl - Cl
State and explain the relationship between the number of bonds, bond
length and bond strength.
The comparison should include bond lengths and bond strengths of
two carbon atoms joined by a single, double and triple bonds
the carbon atom and the two oxygen atoms in the carboxyl group of a
carboxylic acid.
4.2.4
Electron density
(blue shading) is
highest around the
nuclei.


A covalent bond is the electrostatic attraction between a pair
of electrons and positively charged nuclei.
Covalent bonds form between two non-metals that share
their electrons in order to fill the outer energy levels.
2 H atoms
H2 molecule with
1 covalent bond
video
•.
MOLECULES
•
•
A molecule is a discrete group of non-metals atoms
covalently bonded to one another.
The shared electron pair between two bonded atoms
are the bonding electrons.


A scientist called Gilbert Lewis proposed the octet rule: that atoms
change their number of electrons to acquire the stable electron
structure of a noble gas which has 8 valence electrons.
Lewis structure diagrams are a diagrammatical way of showing
valence electrons around elements and in molecules.
Lewis Structures for the first 18 elements:
Pairs of electrons belonging to one atom are called lone
pairs. These are not involved in bonding.
 The unpaired electrons are the bonding sites. These
unpaired electrons are the ones that will be shared with
other atoms during bonding.




Place the central atom, which is the one with the most bonding
sites (i.e. most unpaired electrons), in the centre and place the
other elements around it.
There can be more that one central atom; they should go side by
side.
Draw the most symmetrical structure possible, leaving no unpaired
electrons and ensuring that the octet rule is followed for each atom
(except H which must have two electrons around it).
XX
H O
X
X
X X
The covalent bonds
are the shared pairs
H
X
X
Lone pairs
The shared pair of electrons is replaced with a line
 The lone pairs are omitted.

H O
H
 Double
covalent bonds
involve two shared pairs
of electrons.
 Triple covalent bonds
have three shared pairs
of electrons.
 The atoms on either side
of the double or triple
bond must each
contribute one of the
electrons forming these
bond types.
1. Add up all the valence electrons.
2. Add one electron for every negative charge or subtract
one electron for every positive charge.
3. Draw the Lewis diagram in square brackets and put the
charge outside.
OHSO42-
F2
PH3
CH3Cl
C3H6 Cl2
N2
CO2
NAMING MOLECULAR COMPOUNDS
1.
2.
3.
4.
5.
6.
Write the name of the most metallic atom first (it will be
furthest left on periodic table).
Write the name of the second element with an ide ending.
Add a Greek prefix to each name indicating how many
atoms are bonded.
Exceptions: Do not add the prefix MONO to the first name.
Drop the ‘a’ or ‘o’ from the prefix if the first letter of the
element is the same.
Do not reduce the subscript numbers to the lowest whole
number
Charges are not used to determine the formula of
molecules.
EXAMPLES
N2O dinitrogen monoxide
NO2 nitrogen dioxide
GREEK PREFIXES –PAGE 10 NH
GREEK PREFIXES GO ON….
YOU TRY
–PAGE 11 NOMENCLATURE HANDBOOK
tetrarsenic decabromide
As4Br10
 SF6
sulfur hexafluoride
 carbon dioxide
CO2
 boron mononitride
BN
 P2Cl5
diphosphorus pentachloride

NAMING SIMPLE ORGANIC MOLECULES-ALKANES

Carbon forms many organic molecules.
It has four valence electrons, so it can form up to
4 bonds.
 Carbon can form single, double or triple bonds
even with itself. Carbon can bond to a range of
other elements, including H, O, N, S, and Cl

Hydrocarbons are one type of organic
molecules that form from Carbon and
Hydrogen.
 Alkanes are a group of hydrocarbons that
have all single bonds. Their name ends in ANE
 Their general formula is CnH2n+2
 C3H8, C4H10 , C6H14

NAMING HYDROCARBONS (MOLECULES WITH H AND C)
each bond has a
different length
 double bonds are
shorter than single bond
 triple bonds are shorter
than double bonds






the amount of energy
required to break a bond
measured in kJ mol¯1
it is different for each bond
type
double bonds are stronger
than single bonds
triple bonds are stronger
than double bonds
ASSIGNED WORK
Worksheet on Lewis Structure Diagrams
 Section 2.3 Exercises; numbers 1-7 on page 50

NETWORK COVALENT SOLIDS
(GIANT COVALENT MOLECULES)
PAGE 51
4.2


COVALENT BONDING CONT’D
4.2.9 Describe and compare the structure and bonding in the three
allotropes of carbon (diamond, graphite, and C60 fullerene).
4.2.10 Describe the structure of and bonding in silicon and silicon
dioxide.
A diamond is a network covalent substance. These substances
are large crystalline solids that have all atoms linked by covalent
bonds to several other atoms in a continuous three dimensional
array (or network).
 Because every atom is held in a rigid position, these are the
hardest substances known.
 These macromolecules include diamond Cn and silicates such as
silicon carbon SiC, and silicon dioxide SiO2 (silica or quartz;
major component of sand and used to make glass; most
abundant oxide in Earth’s crust). Silicon also fits into this group.

Diamond
crystalline carbon-
Cn
Cubic zirconia
crystalline silicon carbon which is also
in bullet-proof vest ceramic plates
SiC
Quartz
SiO2
CARBON BASED
Carbon can form different network solids because the
bonding patterns differ.
 These forms are called allotropes: different crystalline
structures of the same element that different in physical
and chemical properties.

ALLOTROPES OF CARBON- 1. GRAPHITE

Each C is covalently bonded to 3 other carbons, forming hexagons which are in
parallel layers and have 1200 angles. It has a 2D arrangement
The layers are only held together by weak van der Waals’s forces (described
later) which allow the layers to slide past one another; graphite is a soft,
powder.
 Because of the sliding layers, graphite is used as a lubricant and is the “lead” in
pencils. Non-lustrous, and grey.
 Since each carbon is only bonded to three others, there is ONE DELOCALIZED
ELECTRON per atom which can conduct electricity as it moves across the
sheet. Graphite is a good conductor of electricity.

GRAPHITE
ON
PAPER
ALLOTROPES OF CARBON- 2. FULLERENE C60





Discovered in 1985, buckministerfullerenes are named after
architect R. Buckminster Fuller who designed the geodesic
dome of the World Exhibition Center in Montreal.
In “buckyballs” each carbon is bonded in a sphere of 60
carbon atoms consisting of 12 pentagons and 20 hexagons.
Structure is a closed spherical cage in which every C is bonded to
3 others.
Yellow, crystalline soft solid (imagine the ball room at
McDonald’s).
Intermediate surface of the “ball” is conductive, but imagine an
ant running across the surface of basketballs.
Reacts with K to make superconducting crystalline
material. Related forms are used to make nanotubes for the
electronic industry, catalysts and lubricants.
ALLOTROPES OF CARBON- 3. DIAMOND




The carbons (Cn ) are arranged in a 3-D lattice where
all the atoms are held together by 4 strong covalent
bonds at 109.50.
These are the hardest minerals known to man; e.g. a
diamond. Used in tools and machinery for grinding
and cutting glass.
Diamonds form when carbon crystallizes under great
pressure. Lustrous crystal.
Unlike graphite, diamonds have no free electrons and
do not conduct electricity.
diamond
6-sided rings (hexagons)
SILICON BASED
SILICON
Silicon is the most abundant element in the Earth’s
crust after oxygen; in the form of silica (SiO2) in sand
and in the silicate minerals.
 It has metallic and non-metallic properties; is hard
and brittle and a most widely used semi-conductor
(computer chips).
 Silicon like carbon is in Group V and has four valence
electrons.
 Each silicon is bonds to four other Si atoms to form a
tetrahedral arrangement like C in a diamond. BUT Si
–Si bonds are longer and weaker than C-C since the
Si atom is much larger and the atoms can not get as
close.

SILICON DIOXIDE-SIO2





Silicon dioxide, also called silica or quartz, is giant covalent
structure.
It is a tetrahedral structure like a diamond but each Si atom
bonds to 4 oxygen and each oxygen bonded to 2 Si. SiO2
refers to the ratio within the molecule.
Strong, insoluble in water, high melting point.
DOES NOT conduct electricity or heat. All electrons are held
tightly between atoms and not free to move.
Sand and glass are different forms of silica.
Quartz
Property
Ionic Compounds
Metals
Network Covalent
Representative
Unit
Formula Unit
NaCl CaF2
Element Symbol
Hg, Cu, Fe
Element Symbol
Diamond (Cx),SiC, SiO2
Bond Formation
Electrostatic attraction
between oppositely
charged ions; and
exchange of electrons
Cations in a sea
of delocalized electrons
Electrostatic attraction
between the nuclei of an
atom and a pair of
electrons; a sharing of e-
Type of Elements
Metallic and
nonmetallic
Metals and
Metals
Non-metal and nonmetal
State at Room
Temp.
Crystalline solid
Soft to hard solids (Hg is
an exception)
Very hard crystalline
solid
Melting Point
High (usually >3000C)
Low to high
Very high
Electrical
Conductivity
Good conductor as
liquid
or dissolved in water;
free moving ions
Excellent conductor
as a solid or molten,
free moving electrons
Poor conductors; no ions
or electrons
Solubility in water Most - high
Not soluble
Insoluble
Other physical
properties
Shiny, malleable
Ductile, dissolves in
Other metals to form
Hardest substances
known to man.
Hard and brittle
REVIEW


How does the structure of a diamond and graphite
differ?
Diamond
Giant Molecule
Covalent bonds only
Graphite
Covalent bonds and
cross-linkage
Layer structure
Compare and explain the hardness and electrical
conductivity of diamond and graphite
Diamond
Graphite
Non-conductor
Good conductor
No delocalized electrons
Delocalized electrons
hard
soft
Rigid structure
Layers can slide
ASSIGNED WORK
Read page 51-56
 Page 57 Section 2.4 Exercises #1-7

Enzyme
receptor
site
Small sugar molecule
THE SHAPE OF MOLECULES
VSEPR
4.2
COVALENT BONDING CONT’D
4.2.7
Predict the shape and bond angles for species with four,
three, and two negative charge. centres on the central atom
using the valence shell electron pair repulsion theory
(VSEPR).
Know the examples: CH4, NH3, H2O, NH4+, H3O+, BF3, C2H4,
SO2, C2H2, CO2
4.2.2
Describe how the covalent bond is formed as the result of
electron sharing
Dative covalent bonds are required. Examples include CO,
NH4+, H3O+
15 MIN READING

Molecular Beauty- Odd Shapes with Useful Functions
The shape of molecules is critical to the way they
interact; especially in biochemical system.
 “Every medicine you take, odour you smell, or flavour
you taste depends on part or all of one molecule fitting
physically together with another. “ (Silverberg, 2006)
 Odour is based on shape not composition. All of these
four substances smell like moth-balls even though they
have a different composition.

Valence-Shell Electron-Pair Repulsion Theory
 Valence electrons (both bonding pairs and lone pairs)
will arrange themselves around a central atom in such as
way to minimize electrostatic repulsion.
repulsion
The repulsion of bonding pairs and non-bonding pairs is
important.
 This repulsion gives all molecules a three-dimensional shape.

Repulsion
Non-bonding (lone
pair) + bonding
Non-bonding
+ non-bonding
>
non-bonding
+ bonding
>
bonding
+ bonding
1.
2.
3.
4.
Draw the Lewis Dot Structure of the
molecule.
Determine the central atom(s); If there
are two central atoms (multiples of C N
or O) each has a shape around it.
Determine the total number of bonds
and lone pairs surrounding the CENTRAL
atom (these are called charge centres;
lone pairs, single, double and triple
bonds all = 1 charge centre)
 For SL Chem, you have to know
shapes with 2,3, or 4 charge centers.
Match the information to the chart.
NH3 has 4 negative
charge centres
O=C=O
H-C≡O
Each carbon has 2
negative charge
centres or 2 bonds and
0 lone pairs
•
•
•
four single bonds; 0 lone pairs = 4
charge centres
bond angles are 109.5o
CH4 (methane), SiCl4, CH3F
•
•
•
one lone pair and 3 single bonds =
4 charge centres
bond angle are 107o
NH3, NF3, AsCl3, PCl3
Repulsion by lone
pair results in
smaller angle than
109 or 120.
•
two lone pairs and two bonds =
4 charge centres
•
bond angle is even smaller: 104.5o
•
H2O, H2S, SI2
•
•
three bonding pairs around the central atom (3 single
or 2 single + a double bond) = three charge centres
0 lone pairs
• bond angles are 120o
• BF3, SO3, CH2O,
(3600/3)
•
2 bonds around the central atom; 0 lone pairs= 2
charge centres
•
combination of single, double, or triple bonds
•
bond angle is 180o
(the bond angle is the angle between the atoms bonded to the
central atom)
•
CO2, HCN, , H2, O2, N2
SPECIAL CASES OF COVALENT
BONDING TO KNOW –PAGE 44
A dative covalent bond is formed by one atom donating BOTH
electrons (shown by arrow in structural diagram).
 These bonds still follow the octet rule
 NH4+ forms when an H+ bonds to the lone pair of N in an
ammonia molecule.
 1N(5) + 3H(3) + 1H+ (0) = 8e- NOT 9e+
H
**
**
N
+
N
H
H
H
H
H
H
H
NH3
hydrogen ion
(1 lone pair)
(no electrons)


The shape is TETRAHEDRAL (like CH4) with 109.50 angles.
dative
bond
Forms when an H+ bonds to the lone pair of O in an
water molecule; another dative bond.
 The resulting H3O+ ion is charged
 The shape is TRIGONAL PYRAMIDAL (like NH3) with bond
angles of 1070.

The C atom shares two electron pairs with O.
 O makes an additional dative bond so that there is a
triple bond (≡) between C and O.
 The shape is LINEAR since there are only two atoms
involved.





S has 6 valence electrons. Oxygen also has 6 valence
and needs to make two bonds to fill its valence shell.
Two double bonds are made with oxygen and two
lone pairs are left on sulfur. This makes 10 electrons
around the S (an exception to the octet rule).
SO2 is a bent molecule but there are 3 charge centres.
(See summary sheet)
Because of the lone pair on the sulfur, the angle is
1170 .
Note SO2 can also be considered dative
 (see handout on dative bonds and exceptions to the octet
rule)


Other exceptions include BeCl2 (16 e-) and BF3 (24 e-)
ASSIGNED WORK
Review shapes and angles; read pages 40-46 of text
 VSEPR Lab
 Complete questions 1-11 on page 50, Exercises 2.3 of
text.

REVIEW
1.
2.
3.
4.
5.
6.
7.
What does VSEPR stand for?
Valence Shell Electron Pair Repulsion (theory).
What causes a molecule to have a specific shape?
Both bonding pairs and lone pairs will arrange themselves around a
central atom in such a way to minimize electrostatic repulsion.
How is a Lewis dot diagram different from a structural diagram?
Lewis dot diagrams show all electrons as x’s or dots, including lone
pairs; a structural diagram only shows the bonding pairs as lines.
Cl has 17 electrons in total. Why does the Lewis dot diagram of Cl
only have 7 electrons?
Lewis dot diagrams only show the valence electrons.
How many lone pairs in this molecule?
9 pairs
How many around the central atom?
0 pair
Is the octet rule being followed?
NO. B is short 1 pair of electrons.
REVIEW
SHAPE
ANGLE
linear
1800
trigonal planar
1200
bent (SO2)
1170
tetrahedral
109.5
trigonal pyramidal (1 L.P.)
1070
Bent (2 L.P.)
104.50
SAMPLE QUESTION
(i) State the shape of the electron distribution around the
oxygen atom in the water molecule and state the shape
of the molecule.
 The electron distribution is tetrahedral.
The shape of the molecule is bent.
(ii) State and explain the value of the HOH bond angle.
 104.5°
Lone pairs repel each other more than bonding pairs;
(INCORRECT to say repulsion of atoms)
PRACTICE- WHAT IS THE SHAPE OF THESE MOLECULES?
QUIZLET
BOND POLARITY AND
ELECTRONEGATIVITY
PAGE 46
4.2



COVALENT BONDING CONT’D
4.2.5 Predict whether a compound of two elements would be
covalent from the position of the elements in the periodic table or
from their electronegativity values
4.2.6 Predict the relative polarity of bonds from electronegativity
values.
4.2.8 Predict whether or not a bond is polar from its molecular
shape and its bond polarities.
Electronegativity is the measure of the ability to attract
shared electrons to itself when is chemically combined
with another atom
 It is measured from 1-4 using the Pauling scale

C
E1
Cl
E2
Electronegativity
2.5
Electronegativity
3.0
Cl has a higher electronegativity
and pulls the shared electrons
closer to itself
The most electronegative
element is fluorine, while
the least electronegative
element is francium
 Electronegativity
increases from left → right on the
periodic table
 It decreases going down a group
 Metals have low electronegativities
 Non-metals have high electronegativities
 F is the most electronegative element
IONIC AND COVALENT CHARACTER CAN BE PREDICTED
FROM TWO FACTORS
In order to form a ionic compound, the elements
reacting together must have very different tendencies to
lose or gain electrons.
 Two ways to predict this:

Position in the periodic table
1.


metal + non-metal = ionic
non-metal + non-metal = covalent
Difference in electronegativities of two atoms in bond.
2.
1.
2.
3.
If the difference in electronegativities is 1.8 or more the bond is
ionic.
If the difference is 0, the bond is non-polar covalent
Between 0 and 1.8 there is a continuum of varying polar covalent
bonding.

Example problem. Determine the type of bond in KF.
ΔEN = ENF – ENK
= 4.0 – 0.8
= 3.2
Since this is > 1.8 KF is an
ionic bond.
A bond dipole (polar covalent bond or permanent
dipole) is when the bond (between two atoms) is
polarized with permanent positive and negative areas.
 The atom with the higher electronegativity has the
stronger electron attraction and has a slightly negative
charge (written as δ‾ ) since the electrons are pulled
closer to it. δ = Greek letter ‘d’ delta.
 The other atom has a slightly positive charge (written as
δ+).
2.1
3.0

δ+
δ‾
H
Cl
δδ+
δ+
PRACTICE- WORKSHEET ON ELECTRONEGATIVITIES
HCl consists of only 1 bond
which is polar.
 Therefore, the HCl molecule is
polar; it has a positive and
negative end.
 Two HCl molecules will be
attracted to each other at their
oppositely charged ends.
 This intermolecular attraction
between molecules is important
because it is the force that
allows molecules to form solids
and liquids to form.

H
Cl
H
δ-
δ+
Cl
Polar bonds such as the H – Be
do not always make polar
molecules.
 BeH2 is composed of two POLAR
B-H bonds.
 BUT it is a NONPOLAR molecule
because the individual dipoles
of the B- H bonds are directed in
opposite directions and cancel
each other.

δ-
δ2δ+




pyramidal and bent are
always polar since they
contain central atoms with
lone pairs; N and O family)
asymmetrical atoms are
polar
symmetrical atoms are nonpolar
tetrahedral, trigonal planar
and linear are non-polar if
symmetrical
Non-polar molecules:

CxRy, where C is carbon
and R is any other nonmetal



CH4, C2F4
Rx, any molecular
element

Polar molecules:
F2, P4, S8
RS where R and S
represent two nonmetals


CxRySz, where C is carbon


HCl, IBr
CH3Cl, C2H4Cl2
NRx or ORx,where N is
nitrogen and O is oxygen

CH3OH, NH3 (not CO2)
INTERACTIVE ANIMATION FOR MOLECULE
POLARITY
YOU TRY.
H
H
A= A
C
H
NONPOLAR
NONPOLAR
H
O
X
X
X
POLAR
X–A–Y
POLAR
X
C = C
X
X
NONPOLAR
Y
X
X
POLAR
X
C
Y
X
C = C
X
X–A–X
NONPOLAR
X
POLAR
YOU TRY- WHAT IS THE SHAPE AND THE POLARITY?

H2O
BENT POLAR

SiCl4
TETRAHEDRAL NONPOLAR

PH3
TRIGONAL PYRAMIDAL POLAR

CH2O
TRIGONAL PLANAR POLAR

CO2
LINEAR NONPOLAR
YOU TRY.
X
X
X
X
Y
Z
C
Y
Y
**
N
X
X
C=C
NONPOLAR
NH3
POLAR
X
C=C
Y
POLAR
POLAR
CH4 NONPOLAR
POLAR
CH3F POLAR
ASSIGNED WORK
Worksheet on polar molecules
 Complete numbers 8-14 on pages 50-51

INTERMOLECULAR
Page 57
BONDING
4.2
COVALENT BONDING CONT’D
4.3 Intermolecular forces
 4.3.1 Describe the types of intermolecular forces (attractions
between molecules that have temporary dipoles, permanent
dipoles or hydrogen bonding) and explain how they arise from the
structural features of molecules.
The term van der Waals’ forces can be used to describe the
interaction between non-polar molecules.
 4.3.2 Describe and explain how intermolecular forces affect the
boiling points of substances.
 The presence of hydrogen bonding can be illustrated by comparing:
HF and HCl, H2O and H2S, NH3 and PH3, CH3OCH3 and CH3CH3OH,
CH3CH2CH3 and CH3CH2OH
Chemical Bonds -WITHIN molecules or between atoms
Ionic
Covalent
Metallic
Network
Physical Bonds
Covalent
Intermolecular:
BETWEEN Molecules
Hydrogen
Bonding
Dispersion
Dipole
Property
Ionic Compounds
Network Covalent
Molecular
Compounds
Representative
Unit
Formula Unit
NaCl CaF2
Element Symbol
Diamond (Cx),SiC, SiO2
Molecule
CH4, F2, H2O
Bond Formation
Electrostatic attraction
between oppositely
charged ions.
Electrostatic attraction
between the nuclei of
an atom and a pair of
electrons.
Intermolecular forces
between molecules (van
der Waals, diplole,
hydrogen))
Type of Elements
Metal and
nonmetal
Non-metal and nonmetal
Non-metal
State at Room
Temp.
Crystalline solid
Very hard crystalline
solid
Solid, liquid or gas
Melting Point
High (usually >3000C)
Very high
Low (usually < 3000C)
Electrical
Conductivity
Good conductor as
liquid
or dissolved in water
Poor conductors
Poor to non-conducting
Solubility in water
Most - high
Insoluble
Varies depending on
polarity
Organic Compounds: fuels (methane CH4, propane C3H8,
butane C4H10 O2, C2H6 (ethane), C2H4 (ethene), and C2H2
(ethyne)
 alcohols (CH3OH, C2H5OH, C3H7OH)
 H2O, HCN, CO2
 O2, N2
 halogens: F2(g) Cl2(g) Br2(l) I2(s)
 Ammonia: NH3

•When you boil water, you break apart the water molecules.
•You DO NOT break the water into hydrogen and oxygen gas.
•You are not breaking the H-O covalent bonds, but the
intermolecular forces whose strength determines if
molecular substances are solids, liquids, or gases.
•These are very weak forces compared to those
involved in ionic, metallic or covalent bonding.
1.
2.
3.
4.
van der Waals forces (London Dispersion)
Dipole-dipole
Ion-dipole (we will only mention this with dissolving)
Hydrogen bonding
Hydrogen
bonding > Dipoledipole
>
Dispersion



Non-polar molecules(Cl2) have no permanent separation of
charges within their bonds because the shared electrons are
pulled equally.
As the cloud of electrons moves, the density may be at any
moment greater in one area than another.
This means that the molecules become slightly polar for brief
instances on a continuous basis; i.e., they have temporary
dipoles which may induce a dipole in a neighbour.
These temporary dipoles allow for a weak electrostatic
attraction briefly.
 The dipoles cancel each other out over time, but they
are strong enough to hold non-polar molecules (like
waxes) together as soft solids and liquids.







Dispersion forces are present in ALL molecules but may not be the
main forces holding the substance together.
They are the weakest of intermolecular bonds
They are important in NON-POLAR molecules because they are the
main force holding them together.
The more electrons in the molecule, the stronger the dispersion
force.
# electrons ↑ dispersion force ↑ melting/boiling points ↑
If the shape of the molecule is complex rather than linear the
molecules can not get close; dispersion force is weak and b.p.

A similar pattern is seen with the noble gases
Element
# electrons
He
Ne
Ar
Kr
Xe
Rn
2
10
18
36
54
86
Boiling
point (0C)
-269
-246
-286
-152
-107
-62
How many electrons do both Kr and HBr have? 36
 These species are isoelectronic (i.e. they have the same
number of electrons).
 Would you expect them to have the same boiling point?
 Kr -b.p. -152 0C
 HBr -b.p. -67 0C
 Some other factor is affecting intermolecular bonding in
addition to the dispersion forces.




The electrostatic attraction between oppositely
charged ends of POLAR molecules
(electronegativity differences).
The positive end of one molecule attracts the
negative end of another molecule.
These are permanent dipole attractions.




If you have two molecules with the same number of electrons, but
one is polar, the polar one will have a higher m.p. and b.p.
This is why the b.p. of HBr is higher than Kr.
However, the dispersion forces are still important for some
molecules. HCl is more polar (∆EN = 0.9) than HBr (∆EN = 0.7) but
HCl has a lower b.p. suggesting that HBr has stronger bonds. WHY?
What would you expect the boiling point of HF to be?
BOILING POINTS FOR HYDROGEN COMPOUNDS OF GROUP 4, 5, 6 AND 7 (PAGE 61)
The electrostatic attraction between the nucleus of a
hydrogen atom bonded to a highly electronegative
element (H-F, H-O or H-N) and the lone pair of a F O
or N of the nearby molecule.
 F O and N are electronegative atoms which pull the
electrons away from H leaving it with a positive
charge.


F
CH3F can not form H-bonds because
the H is bonded to C not F.
H
C
H

But ammonia does form H-bonds.
H
**
N
H
H

It is also very soluble
in water; WHY?
H



All H-bonded compounds are polar (HF, NH3, H2O)
Important in biological systems
The strongest of intermolecular bonds.


All molecules with an OH group are alcohols.
The OH of the alcohol form a H-bond with the water
molecule. Therefore alcohols are soluble in water (to
different degrees).
H-F Bond The only example is hydrogen fluoride
O-H Bond Example: Methanol CH3OH
N-H bond Example: Urea (CO(NH2)2

1.
2.
3.
4.
Responsible for many of the special properties of
water:
High surface tension
Solid water is less dense than the liquid
High specific heat capacity; takes a lot of energy to
heat it compared to the same amount of a metal.
High melting point
1. Surface tension
meniscus
2. Solid is less dense that liquid
2. Structure of Ice- When liquid water changes to ice.
Hydrogen bonding between the molecules is directed by
the v-shape of the water molecule and leaves hexagonal
holes making ice less dense. The H-bonds stabilize the
structure.
3. Specific heat capacity (amount of heat needed to raise
the temperature of 1g of a substance 1 degree). Water
has a high specific heat capacity and is slow to warm
up and cool down (unlike metals, for example).
2.6 PHYSICAL PROPERTIES
Page 63
4.5
PHYSICAL PROPERTIES
4.5.1 Compare and explain the properties of substances
resulting from different types of bonding
Examples should include melting and boiling
points, volatility, electrical conductivity and
solubility in non-polar and polar solvents
REVIEW OF PHYSICAL AND CHEMICAL PROPERTIES
 Physical
property: or condition of a substance
that can be observed or measured without
changing the substance’s composition.
 Chemical property: how a substance reacts with
other substances

Chemical change involves the breaking of bonds that
hold the atoms of the different elements together, and
the formation of new bonds and a new substance(s).
EXAMPLES OF PHYSICAL AND CHEMICAL PROPERTIES
Physical Properties
Qualitative
(descriptive)
Quantitative
(measurable)
Physical state at SATP*
Colour
Odour
Crystal shape
Malleability
Ductility
brittleness
Melting/freezing point
Boiling/condensing
point
Density
Solubility
Electrical conductivity
Thermal conductivity
Chemical Properties
Reactivity with water
Reactivity with air
Reactivity with pure oxygen
Reactivity with acids
Combustibility (flammability)
toxicity
*SATP (standard ambient temperature and pressure) = standard
conditions of 250C and 101.3kPa (kiloPascals) of pressure


Relative boiling points of substances can be predicted knowing
the types of bonds involved; the stronger the bond the higher
the melting and boiling point.
Bonds types can be ranked from strongest to weakest:
Network covalent > ionic > metallic > hydrogen > dipole-dipole > van der Waals
(varies)





Polarity also affects solubility. “LIKE DISSOLVES LIKE”
Water is a polar molecule. Only polar substances will
dissolve in water. NaCl will dissolve in H2O.
Ion-dipole forces are between an ion and a polar solvent
like water.
Non-polar substance dissolve in non-polar solvents (oils,
gases).
Wax (C25H52 (s) )will dissolve in cyclohexane (C6H12 (l) ) but
salt will not.
Volatility is how easy a substance turns into a gas or
vapour.
 Weak bond enable atoms or molecules to escape easily
from the surface.
 Molecular substances are volatile, while ionic
compounds, metallic compounds (except Hg) and
network covalent are not volatile.

ASSIGNED WORK



Handout Intermolecular bonding
Section 2.5 page 62; questions 1-4 (these are good practice)
Worksheet 2.5 Intermolecular Forces (for extra practice)
Physical Properties
 Section 2.6 page 65; questions 1-9
 Worksheet 2.6; Physical Properties (for extra practice)



Chapter Review in text
Handout Bonding Overview –The inside 2 sheets are the ones
to focus on (fill in the blank)
Worksheet 2.7 Bonding Summary (extra practice)
YOU HAVE
SURVIVED….
BONDING
IB TOPIC 4
Chapter 3