
The nucleus is made up of protons (p+) and
neutrons (no) and is surrounded by rings of
orbiting electrons (e-)
Subatomic
particle
Relative mass
Relative charge
Proton
1
+1
Electron
0
-1
Neutron
1
0

Also called isotope notation
A
Z
X
X = element symbol
A = atomic mass = # protons + # neutrons
Z = atomic number = # protons

What is the difference between these two
atoms?
Introduction: BONDING



Chemical reactions involve the formation of
new products
Bonds between atoms or ions in the reactants
must be BROKEN (this requires energy, and is
therefore an ENDOTHERMIC process)
Bonds are then FORMED between atoms or
ions to make the products of the reaction
(this releases energy, and is said to be an
ENDOTHERMIC process)
Intramolecular vs Intermolecular Forces



The bonds that are present between atoms within a
molecule are due to INTRAmolecular forces of
attraction – relatively STRONG forces of attraction
that hold the molecule together
Forces of attraction between molecules,
INTERmolecular forces of attraction, are what keeps
molecules together as a solid, or liquid;
INTERmolecular forces of attraction must be
overcome if a substance is to change its physical
state from solid to liquid to gas: The strength of
these intermolecular forces is affected by the type of
bonding within the molecule
o
We will begin with a study of the types
of bonding found in compounds, and
then look at the characteristics of
these bonds and how it affects the
PHYSICAL PROPERTIES (like solubility,
melting point, boiling point,
conductivity and hardness) of different
compounds




Ions are elements that have gained or lost
electrons
Ions are commonly found dissolved in water,
such as in the cytoplasm or plasma of the
blood
Elements in the same family tend to form the
same type of ion
(e.g.: Na+, Li+, K+, Rb+)
Some important ions are Ca2+ (used for muscle
contraction), Na+ and K+ (nerve and muscle
function), Fe2+ and Fe3+(in hemoglobin) and H+
(required for synthesis of ATP)



Electrons orbit the nucleus of an atom at a
great distance compared to the size of the
particles
Analogy: If an apple represented the size of
an atom’s nucleus and it was placed at the
center of the earth’s core, the valence
electrons would be orbiting close to the
surface of the earth’s crust
The valence electrons therefore are the part
of the atoms that interact in chemical
reactions to form compounds




Ionic Compounds form when electrons are
transferred from one atom to another to form ions
with complete outer shells of electrons (ie the same
stable electron configuration as the inert gases.)
Ionic BONDS form between a metal and a nonmetal
eg NaCl
Metal tends to lose electrons which are transferred
to the nonmetal
Metals form a cation (+) and nonmetals form an
anion (-)
Formation of NaCl
Formation of MgF2
Ionic Compounds



Positive and negative
ions are held together in a
strong lattice framework
through the strong
electrostatic attraction of
oppositely charged ions
Ionic compounds have
high melting points and
are solids at room
temperature
http://wps.prenhall.com/wps/media/ob
jects


These result in a lattice of ions rather than
individual molecules, so we refer to MgF2 and
NaCl as formula units, not molecules.
Properties of ionic substances as a result of
the strong, rigid lattice that can dissociate
when dissolved in water:
◦
◦
◦
◦
◦
Crystalline solids at room temperature
Hard and brittle
High melting and boiling points
Conduct electricity when in liquid form
Most are soluble in water
Ionic Compounds are crystalline:
Foodsubs.com
Docbrown.info





Dissolving of NaCl in water:
http://www.youtube.com/watch?v=EBfGcTAJF
4o&feature=related
http://www.youtube.com/watch?v=dr4sFNzU
VzI&feature=related
NaCl solutions conduct electricity
http://www.youtube.com/watch?v=aELPrWzix
eU&feature=related


The chemical formula of simple ionic
compounds can easily be determined if one
knows how many electrons an atom needs to
lose or gain to achieve a stable inert gas
electron configuration – this is related to its
group number = # of valence electrons
already present
In an ionic compound, the sum of positive
and negative charges = zero to create a
neutral compound e.g. MgS, Al2O3 CaBr2
Group # indicates charge for
families 1,2,3 and 5,6,7
1
2
3
+1
+2
+3
Li
Be
Na
Mg
Al
4
P
5
6
7
-3
-2
-1
N
O
F
S
Cl
Formulas and Names of Ionic
Compounds

Simple Ionic Compounds: metal + nonmetal

Eg sodium and bromine
Names and Formulas of ionic
compounds continued…….

Calcium and chlorine
Simple Ionic Compounds

To name, state the name of the metal
element first, the nonmetal element second
and change the ending of the nonmetal
element to “ide”. No prefixes are used.

eg NaCl

Eg MgBr2
sodium chloride
magnesium bromide
Words  formula


Use the crossover rule to determine the
formula, using the valence or ionic charge of
each ion. The resulting compound should be
neutral (net charge = zero)
Eg aluminum oxide
eg lithium bromide
For ionic compounds, always reduce the ratio of
metal : nonmetal to lowest terms

Eg
Magnesium and oxygen
Alternatively, use an accounting
method instead of the crossover:

Calcium and oxygen or calcium and chlorine
Multivalent Metal Ions
Iron (II)
Fe2+
FeCl2
iron (II) chloride
Iron (III)
Fe3+
FeCl3
iron (II) chloride
Copper (I)
Cu1+
Cu2O
copper (I) oxide
Copper (II)
Cu2+
CuO
copper (II) oxide
Tin (II)
Sn2+
SnO
Tin(IV)
Sn4+
SnO2
tin (IV) oxide
Lead (II)
Pb2+
PbO
lead (II) oxide
Lead (IV)
Pb4+
PbO2 lead (IV) oxide
tin (II) oxide
Polyatomic ions

Some ions are composed of a number of
nonmetal elements bonded together
covalently, with an overall net charge due to a
surplus or deficit of electrons

Eg NH4+ the ammonium ion

Eg SO42- the sulphate/sulfate ion
Common Polyatomic ions
NH4+
ammonium
SO42-
sulfate
OH-
hydroxide
SO32-
sulfite
NO3-
nitrate
HSO42-
hydrogen
sulfate
NO2-
nitrite
PO43-
phosphate
ClO3-
chlorate
HPO42-
hydrogen
phosphate
CO32HCO32-
Carbonate
Hydrogen
carbonate
H2PO41-
dihydrogen
phosphate
Memory Aid for Polyatomic Ions

“Nick the Camel ate a Clam for Supper in Phoenix”

First letter:

# consonants = # of O’s in formula



the non metal element
# vowels = # of negative charges on the ion
OR
# of hydrogens in the related oxyacid
Example:
Nick

“N” is for nitrogen

3 consonants = 3 O’s

1 vowel = 1 negative charge

or
1 H in the oxyacid:
N
NO3
NO3
–
HNO3(aq)
Now you try…….

Camel tells us:
Nick the Camel gives us the –ate
polyatomic ions, aka the “MOTHER
IONS” (by Mrs. Wheelihan)



When 1 oxygen is removed from the oxyion,
the ending changes from -ate to –ite, BUT
THE CHARGE STAYS THE SAME
EG NO3- nitrate and NO2- nitrite
SO42- sulfate and SO32- sulfite
This idea can be used to name a number of variations
from the “mother ions”
Sometimes a hydrogen ion, H+,
stays attached to a polyatomic ion


This creates a new polyatomic ion with the
prefix “hydrogen” or “dihydrogen”
Eg
hydrogen carbonate ion
Other examples:

PO43HPO42H2PO41-

NaH2PO4


phosphate
hydrogen phosphate
dihydrogen phosphate
sodium dihydrogen phosphate
Peroxide Ion, O22



PEROXIDES contain the peroxide ion, O22- instead
of the oxide ion, O2example, hydrogen peroxide H2O2
Peroxides are named by placing the prefix “per”
before the word oxide in the name of the
compound. When writing the formula of a
peroxide, simply add an additional oxygen atom
to the formula of the normal oxide, representing
the peroxide ion, O22Eg lithium peroxide Li2O2




Form between two nonmetals
Electrons are shared rather than transferred
Macromolecules and organic molecules are
covalent molecules held together with
covalent bonds.
Examples: lipids, carbohydrates, proteins
and nucleic acids.
What holds a covalent bond together?
 Nuclei repel each other, but are both
attracted by the pair of negative electrons
being shared.
Formation of H2
Formation of NH3
Other examples, with their Lewis diagrams:
Formation of O2 – a double bond
“Lone pairs” and “Shared pairs”




Non-bonded pairs are called “LONE PAIRS”
Pairs of electrons that are shared between
atoms are called: ”BONDING PAIRS”
BOTH LONE PAIRS AND BONDING PAIRS MUST
BE SHOWN ON YOUR LEWIS DIAGRAM
ie ALL VALENCE ELECTRONS must be shown
Bonding pairs and Lone Pairs
Alevelchem.com
Lewis Structures of Covalent Molecules


Covalent molecules can be represented in a
number of different ways. In Lewis structures,
ONLY THE VALENCE ELECTRONS ARE SHOWN.
In Lewis structures, the outside electrons are
shown with dots and covalent bonds are
usually shown by bars, but there are several
ways of drawing Lewis structures (see Figures
5 to 9, Neuss (2007) pages 64, 65)
There are a number of different ways to represent
Lewis structures of covalent molecules:
These can also be drawn with lines or crosses to represent lone pairs and bonding pairs
of electrons
. Covalent Bonds usually involve the sharing of an
electron pair consisting of 1 electron from each
molecule involved in the bond
Chemteam.info
Coordinate or Dative Covalent
Bonds


Sometimes the “bonding pair” of electrons
both come from the same element. This type
of bond is called a “coordinate or dative
covalent bond”.
Eg ammonia
Chewtychem.wiki.edu
Other examples of dative covalent bonds (also called coordinate bonds)
are found on page 65 of Neuss (2007) and include carbon monoxide,
CO(g); the hydronium ion, H3O+; and aluminum chloride dimer
Single Covalent Bond

When 1 pair of electrons is shared between
the same two atoms, a SINGLE COVALENT
BOND is formed
Double Covalent Bonds


When 2 pairs of electrons are shared between
the same two atoms, a DOUBLE COVALENT
BOND is formed as found in CO2 (g), carbon
dioxide or carbon (IV) oxide
Carbon dioxide CO2
http://dbhs.wvusd.k12.ca.us
Triple Covalent Bonds


When 3 pairs of electrons are shared between
atoms, a triple bond is formed
Eg, nitrogen gas, N2
http://www.webchem.net/notes/chemical_bonding/covalent_bonding.htm
BOND LENGTH and BOND STRENGTH


The more pairs of electrons shared between
atoms, the stronger and shorter the bond
Therefore single bonds tend to be longer
and weaker than double bonds and triple
bonds are shorter and stronger than
corresponding double bonds, as shown in
Table 1 below
Table 1: Bond length and bond strength (Neuss 2007 p. 66)
Example
O-O
Cl-Cl
N-N
C-C
H-H
0=0
N=N
C=C
C≡C
N≡N
Length (nm)
Strength (kJ mol-1)
If single, double, and triple C to C bonds
are compared a distinct pattern emerges:
Bond
Bond
strength
kJ/mol-1
Bond
length
nm
C-C
348
0.154
C=C
612
0.134
C=C
837
0.120
Describe the trend in bond strength and bond length observed in this table:
Naming Covalent Molecules

DIATOMIC MOLECULES – Gaseous elements

HOBrFINCl stands for H2, O2, Br2, F2, I2, N2, Cl2


ALL are diatomic molecules, whose name is
the same as the element name
Eg chlorine Cl2
10. Molecular Elements






Some elements combine with themselves to form
molecules and this is the way that they are found in
nature.
H2(g) hydrogen gas
P4(s) phosphorus
F2(g) fluorine gas
N2(g) nitrogen gas
S8(s) sulfur
Cl2(g) chlorine gas
O2(g) oxygen gas
Br2(l) bromine
I2(s) iodine
“HOBrFINCl”
Binary Acids and their Gases


There are 5 common binary acids, that form
when the corresponding gas dissolves in
water to form an acidic solution. The gas is
given a different name from the acid so that
we can tell them apart and we use subscripts
indicating their state with the formula to tell
them apart.
The binary acids consist of HYDROGEN and a
second NON-METAL. Their acids end with the
suffix -ic.
Hydrogen Halides and their
corresponding Binary Acids
Hydrogen Halide Gases (g)
Binary Acids (aq)
HF(g)
Hydrogen
fluoride
HF(aq)
Hydrofluoric
acid
HCl(g)
Hydrogen
chloride
HCl(aq)
Hydrochloric
acid
HBr(g)
Hydrogen
bromide
HBr(aq)
Hydrobromic
acid
HI(g)
Hydrogen
iodide
HI(aq)
Hydroiodic acid
H2S(g)
Dihydrogen
sulfide
H2S(aq)
Hydrosulfuric
acid
General Rules for Naming Covalent Molecules




Write the LEAST ELECTRONEGATIVE element
first in the name and in the formula.
Naming Method #1: Common method uses
the Prefix system ( also called Classical) the
1=mono
2=di 3=tri
4=tetra
5=penta
6= hexa
7= hepta
e.g. CCl4
carbon tetrachloride
IUPAC and covalent molecules
Naming Method #2: IUPAC system uses
Roman Numeral to indicate the oxidation
number of first element. Refer to text’s
periodic table for possible oxidation numbers
of nonmetals.

CCl4

carbon (IV) chloride
Electronegativity




Linus Pauling developed the concept of
electronegativity (En)
It is a measure of how strongly an atom
attracts electrons to itself in a covalent bond
Fluorine has the highest En value, and Pauling
assigned it an arbitrary value of 4.1
Elements to the left and below fluorine have
decreasing En values
Electronegativity Trends
Ionic or Covalent? Find ΔEn




In bond formation, it is useful to look at the
electronegativity difference (ΔEn)
When Pauling looked at a range of bonds and
their ΔEn values, a pattern was noticed
Bonds with an ΔEn greater than 1.7 tended to
exhibit ionic characteristics
Bonds with an ΔEn below 1.7 tended to
exhibit covalent characteristics
HBr
En (hydrogen) = 2.1
En (bromine) = 2.8
ΔEn = 2.8 – 2.1 = 0.7
HBr is polar covalent
LiF
En (lithium) = 1.0
En (fluorine) = 4.1
ΔEn = 4.1 – 1.0
= 3.1
LiF is ionic
Electronegativity differences between 0.4 and 1.7 mean the
ELECTRON PAIR WILL NOT BE SHARED EQUALLY, BUT WILL BE
closer to the nuclei of the more electronegative element. This
creates a POLAR COVALENT BOND.
The O-H bonds in water are polar
covalent




There are two types of covalent bonds
Atoms that have the same En will have an ΔEn
of zero.
These atoms will attract the shared electrons
equally, and so the distribution of electrons
is uniform
These are nonpolar
covalent bonds


Covalent bonds that have two different
elements will have different En values and so
the electron distribution will be non-uniform
These bonds are called polar covalent, since
one end of the bond will be slightly
electronegative (δ-)since the electrons are
attracted more to the atom at that end



Valence Shell Electron Pair Repulsion theory
(VSEPR) allows us to predict the 3-D shape of
a molecule
VSEPR theory states that bond pairs of
electrons repel one another, and lone pairs of
electrons take up more space than bond pairs
There are four basic shapes which are
common in organic molecules:




Linear
Bent or V-shaped
Tetrahedral
Pyramidal

Linear
or

Bent

Pyramidal

Tetrahedral


How do we determine if a molecule is polar or
nonpolar?
A polar molecule has an uneven distribution
of electrons. This occurs when
◦ There is at least one polar bond
◦ The shape of the molecule is asymmetrical
◦ Or the shape is symmetrical but the atoms
surrounding a central atom have different En values
Methane: CH4
VSEPR diagram:
Polar bonds? Yes
Overall dipole? No
Methane is: NON-POLAR
Ammonia: NH3
VSEPR diagram:
Polar bonds? Yes
Overall dipole? Yes
Ammonia is: POLAR
Water: H2O
VSEPR diagram:
Polar bonds? Yes
Overall dipole? Yes
Water is: POLAR
Carbon dioxide: CO2
VSEPR diagram:
Polar bonds? Yes
Overall dipole? No
Carbon dioxide is: NON-POLAR





The particle theory states that there are
forces between particles, and the forces
increase as the particles get closer.
These are the intermolecular forces
Compared to covalent and ionic bonds, they
are very weak – but when there are many,
they add up to a significant force
Collectively they are called van der Waals
forces, but there are three different forces.
These forces have an effect on the boiling
point and the solubility of substances.
vanderWaal’s forces



vanderWaal’s forces of attraction occur when
the protons in one atom or molecule attract
the electrons in a neighbouring atom or
molecule.
Since all particles have protons and electrons,
all substances have vanderWaal’s forces
Larger molecules have more protons and
electrons, and so have stronger vanderWaal’s
forces.
vanderWaal’s forces of nonpolar
hydrocarbons affect boiling points:
When comparing
the boiling points
of hydrocarbons
(non-polar
molecules), we
see that the
boiling point
increases as the
number of carbons
increases.
Why is this?



Occurs only in polar molecules
that have hydrogen and at least
one of the following atoms: N, O
or F.
These highly electronegative
atoms have lone pairs of
electrons which are attracted to
the hydrogen atoms in
neighbouring molecules.
These hydrogen atoms are
essentially a proton


Polar substances have a slightly
electronegative end and a slightly
electropositive end.
Dipole-dipole forces occur when oppositely
charged poles momentarily attract one
another



Water is not an organic molecule but is
essential for life on this planet
All cells are surrounded inside and out with
water – anything that interacts with a cell
must first be dissolved in water
Physical properties:
◦
◦
◦
◦

colourless and transparent
liquid at room temperature
density = 1.0 g/mL
m.p. = 0℃
b.p = 100℃
water has LD, D-D forces, and H-bonding
Water has cohesive properties – the high
number of intermolecular forces causes water
molecules to ‘stick’ together
Examples:

◦ surface tension – beading of water
◦ water striders – too light to break surface tension
◦ transpiration in plants – transport in xylem tubes
Water has adhesive properties – it’s polar
nature causes it to stick to other substances
Examples:

◦ capillary action – water ‘climbs’ up small diameter
tubes, or ‘bleeds’ through the microscopic pores
and channels in paper or other porous substances
◦ this is due to the hydrogen bonding interactions
between the water and the surface of the tube
(either SiO2 or the cellulose tubes of paper)
◦ This helps to explain the meniscus inside a tube





Water has outstanding solvent properties
Used to be called the ‘universal solvent’, but this is
not a good name, since not everything dissolves in
water
The polar nature of water allows any other polar
substance or any charged particle to dissolve easily
The δ- will attract the δ+ end of solutes, and this
attraction will remain once the solute is dissolved.
The same is true for ionic substances – the cation
will be attracted to the δ- end of water, and the
anion will be attracted to the δ+ of water.
Special Properties of water


Water’s density decreases as it changes from
liquid to solid.
This is because the distance between molecules
in a crystal lattice (as ice) on average further
than when in a liquid.