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Chemistry NCEA L2
2.4 Bonding, Structure and Energy
1
Achievement Criteria
Bonding, structure and energy changes are limited to:
Lewis structures, shape and polarity of simple molecules. Simple molecules have no more
than four electron pairs about any atom (including multiple-bonded species)
Intermolecular forces (the distinction between the different types of intermolecular forces
is not required)
Ionic, covalent and metallic bonding
Molecular, ionic, metallic and covalent network substances
Properties are limited to hardness (including malleability and ductility), electrical
conductivity, melting and boiling points and solubility.
Exothermic and endothermic reactions including energy changes associated with the
making and breaking of chemical bonds, differing amounts of substances and changes of
state. These may involve calculations.
2
Introduction
Chemistry is the study of matter and energy and the interaction between them.
The elements are the building blocks of all types of matter in the
universe. Each element is made up of only one type of atom, each
with its specific number of protons known as its atomic number.
3
Introduction
A large amount of energy is required to break an atom
down into smaller particles. The elements occur in
widely varying quantities on earth. The ten most
abundant elements make up 98% of the mass of earth.
Many elements occur only in traces, and a few elements
are synthetic and highly unstable.
4
Atomic Theory
Atoms are the building blocks of elements.
Scientists and philosophers have guessed that all matter is made up of
building blocks for a very long time. Discovery of the actual structure of the
atom has only been in relatively recent times however.
Atomic Theory
Democritus – Greek philosopher 400 B.C.
The word atom comes from the Greek word atomos meaning indivisible.
Democritus reasoned that you could continue to half a
piece of matter only to a particular point – the smallest
particle that could no longer be divided was known as an
atom. The structure of the atom was unknown at this time.
Extra
for
experts
Atomic Theory
Extra
for
experts
Boyle – English physicist 1661
Robert Boyle investigated the elements,
gold and silver, and found they could not
be changed into earth, air, fire or water,
that previously Aristotle (Greece 350BC)
claimed all matter was made of. He
concluded that gold and silver were true
elements and created a definition for an
element:
An element is any substance that cannot
be broken down into a simpler
substance.
SJ Gaze
Atomic Theory
Extra
for
experts
Lavoisier – French Chemist 1743 - 1794
Antoine Lavoisier used a weighing balance to investigate mass relationships in
chemical reactions. The law of Conservation of Mass is based on his findings.
The total quantity of matter before and after a chemical reaction is the
same.
Atomic Theory
Extra
for
experts
Proust – French Chemist 1799
Joseph Proust investigated individual elements that made up compounds. The
law of Constant Composition is based on his findings.
All pure samples of a given compound contain the same elements combined
in the same proportion by mass.
Extra
for
experts
Atomic Theory
Dalton – English Chemist 1808
John Dalton lived most of his life as a teacher and
public lecturer, but his curiosity of the world
around him led him into research and
experimentation in many areas of Science. One of
his most valuable areas of study was in Chemistry
concerning the atom and elements. Dalton devised
his own notation for elements known at the time.
Atomic Theory
Dalton – English Chemist 1808
John Dalton combined previous ideas about atoms in a
published a
book, called A New System of Chemical Philosophy, adding
his own ideas. A summary of the main ideas from his Atomic Theory are:
All matter is made up of extremely small particles called atoms.
All atoms of one element are identical.
Atoms of one element differ in size and mass from atoms of another
element.
Atoms are indestructible and merely rearrange themselves during chemical
reactions.
Atoms of different elements combine in simple whole number ratios to form
chemical compounds.
Dalton’s Atomic Theory
1. All Matter is made up of atoms
2. Matter can neither be created nor
destroyed
3. Atoms of a particular element are alike
4. Atoms of different elements are different
5. A chemical change (reaction) involves the
union or separation of individual atoms
Atomic Theory
Dalton – English Chemist 1808
Dalton also investigated compounds and came up with the following ideas:
Compounds are composed of molecules
Molecules are composed of atoms in definite proportions
Atomic Theory
Avogadro – Chemist 1811
Amedeo Avogadro also demonstrated that some elements, such as hydrogen
and oxygen, existed in their natural state as more than one atom bonded
together. He called those particles molecules.
A molecule is the smallest particle of a particular substance which can exist
in a free state.
Atomic Theory
Berzelius – Swedish professor 1813
JÖns Berzelius developed a system of chemical
symbols to represent elements. Dalton had
devised little pictures but these were difficult to
draw. He used letters of the alphabet for the
latin name of each element, and a further letter
if there was more than one element that started
with the same name. Berzelius’s system is now
used universally by scientists.
The first letter is always a capital letter
The second (if there is a second) is always a
lower case letter
Atomic Theory
Elements consist of only one type of atom.
Each element can be represented by a chemical symbol.
Most symbols are one or two letters, formed
from the name of the element,
e.g. Hydrogen H, or Helium He.
The first letter of the symbol is always a
capital letter. Any other letters are lower
case. e.g. Helium is He not HE
If the symbols are not based on a elements
English name then it is most likely to be
based on it’s Latin name, the original
language of Science.
Atomic Theory
Extra
for
experts
JJ Thompson – English Physicist 1897
Thompson used cathode ray tubes to discover the presence of negatively
charged particles 1/2000 the mass of atoms. These particles, called electrons,
were found in all atoms. He developed the ‘plum pudding’ model to explain
that the negative electrons must be balanced by positive particles to produce
a neutral atom.
Atomic Theory
Rutherford – New Zealand Physicist 1911
Ernest Rutherford discovered the true structure of atoms
by using alpha particles (helium nuclei) fired at gold foil.
A large number of these particles travelled straight
through and he surmised:
Almost all of the space occupied by an atom must be empty
From the small number of particles that bounced directly back he surmised:
The centre of an atom must be very dense and small
From the way the positive alpha particles were deflected is surmised:
The centre of an atom must contain a region with a very great positive
charge
Atomic Theory
Extra
for
experts
Rutherford – New Zealand Physicist 1911
Results from Rutherford’s gold foil experiment could not be explained by the ‘plum
pudding’ model of Thompson so Instead, in 1911, Rutherford proposed a new model of
the atom in which all of the positive charge is condensed into a tiny, massive nucleus
about ten thousand times smaller than the entire atom. Rutherford explained the much
lighter electrons circulated outside the nucleus.
This was a revolution
in the ideas of atoms
as Rutherford’s model
implied that matter
consisted almost
entirely of empty
space.
Atomic Theory
Ernest Rutherford named the positive particles protons found in the dense centre
nucleus. Because this accounted for only around half the mass in most atoms he
deduced there must be a similar amount of neutral particles - he named these
neutrons.
Electron
-ve
Nucleus
Neutron
N
NP P
N N P
P
P NP N
Proton
+ve
Shell 1
Shell 2
Atom Charges
Protons are positively charged; electrons are negatively charged; neutrons have o
electrical charge.
Atoms have no overall charge because the number of protons = number of electrons.
All matter is made up
of atoms. Atoms
consist of protons,
neutrons and
electrons.
The charges of protons
and electrons are
equal and opposite.
Atom Summary
Subatomic
particle
symbol
Mass
compared to
a proton
charge
location
Proton
p
1
+1
In the
nucleus
Neutron
n
1
0
In the
nucleus
Electron
e
1/1840
-1
Moving
outside the
nucleus
Atomic Theory
Moseley – English physicist 1914
Henry Moseley worked alongside Rutherford in investigating atom structure.
He discovered some elements had different atomic masses even though they
had the same number of protons. He realised elements could be found with
different numbers of neutrons and called these varieties isotopes.
Atoms with the same number of protons but different
numbers of neutrons are called isotopes.
Isotopes
Isotopes of elements occur when atoms have the same atomic number (Z) but
different numbers of neutrons in the nucleus. The numbers of neutrons in an atom
does not affect the way an element behaves chemically, but it does affect the way it
behaves physically.
Isotopes found in nature are generally stable, however radioactive isotopes do exist
such as 238Uranium
24
Atomic and Mass number
The atomic number is unique for each element.
A neutral atom has the same number of electrons as protons.
The periodic table is arranged in order of an elements atomic number.
The mass number is the total number of protons and neutrons together.
Calculating protons, neutrons and electrons
Number of protons:
For an atom or ion = atomic number
Number of electrons:
For an atom = atomic number
For a negative ion = atomic number + charge (- =1, -2 =2 etc)
For a positive ion = atomic number – charge (+ =1, +2 = 2 etc)
Number of neutrons:
For an atom or ion = mass number - atomic number
atom or ion
number of protons
number of electrons number of neutrons
Mg
12
12
12
Mg2+
12
10
12
F
9
9
10
F-
9
10
10
Atomic Theory
Bohr – Danish Physicist 1885 - 1962
Niels Bohr was a physicist who also worked alongside Rutherford. He devised
a new theory to explain how electrons continuously orbit the nucleus and not
gradually lose energy and spiral into it. The main points of his theory were:
Electrons only occupy certain orbits of fixed energy
Electrons that remain in an orbit don’t emit or absorb energy
Electrons can move between orbits. By absorbing energy an
electron can move to an orbit further away at a higher energy
level. By emitting energy (in the form of light) an electron can
move into a lower energy level that is closer to the nucleus.
The electron orbits are called energy levels or shells.
SJ Gaze
Energy Levels in Atoms
Electrons move further away from the nucleus as they gain in energy. Energy
can be provided in the form of heat or light.
28
Quantum physics and energy levels
Extra
for
experts
Why are electrons in shells or energy levels?
It all has to do with quantum physics, where particles, in this case electrons
can only contain particular packages of energy. The more packages of energy
an electron has, the further out shell it sits in from the nucleus. There can be
no half packages or quarter packages; only whole packages, thus the word
quantum meaning quantity.
Classical physics
A turtle sitting on a ramp can have
any height above the ground- and
so, any energy level.
Quantum physics
A turtle sitting on a staircase can
only sit at certain heights, therefore
has only certain energy levels.
The electrons in an atom are arranged in a series of energy levels.
Electrons move or ‘orbit’ around the nucleus in energy levels or shells. The
energy levels further away from the nucleus are able to fit more electrons.
The first energy level is filled first, followed by the second and so on until all
the electrons (the same number of protons in an atom) have been used.
Maximum numbers of electrons in
each energy level are:
>2 in the first EL (nearest the
nucleus)
>8 in the second EL
>8 in the third EL (before the
fourth shell starts to fill)
>18 in the fourth EL
You need to draw the configurations of
the first 20 elements as well as knowing
their names and symbols
Valence electrons
The electrons in the outermost energy level of an atom are known as valence
electrons. These valence electrons are the particles that react with other atoms in a
chemical reaction. The number of valence electrons can determine how an atom will
react with others.
There is a relationship between the period number and the number of
energy levels an atom has.
At this time, the maximum number of
energy levels(or electron orbitals) for
any element is seven.
In the periodic table,
elements have something in
common if they are in the
same row. All of the
elements in a period have
the same number of energy
levels. Every element in the
top row (the first period)
has one energy level for its
electrons) All of the
elements in the second row
(the second period) have
two energy levels for their
electrons. It goes down the
periodic table like that.
Electron configuration
A shorthand way of describing the way electrons are arranged in an atom is called the
electron configuration. The information for the number of electrons is found by an
elements Atomic Number (number of electrons = number of protons in a neutral
atom). Each energy level is filled to its maximum capacity, starting with the lowest
energy level first (energy level number 1 or M shell). The energy level are separated by
a comma. The energy levels are filled until all the electrons are placed.
12
Atomic
number
The total of the electronic configuration
must equal the atomic number in an atom
2, 8, 2
Mg
First EL, second EL, third EL
24
33
Using the Periodic table to write electron configurations
Period number gives
number of energy levels
Last number of group
gives electrons in outer
energy level. i.e. group
17 - 7 electrons in outer
energy level.
Step 1. Ca in period (row 4) so
has 4 energy levels
Ca
2 , 8 ,8 ,2
Step 2. Ca in group 2
so has 2 electrons in
the outside energy
level
Step 3. backfill all energy levels
with 8 electrons (2 in first) and
add commas between each
Extra
for
experts
Atomic Theory
Latest theories in atomic structure.
Atoms consist of a nucleus and surrounding electron cloud. The electrons are obeying
quantum
mechanics; their
The mechanics
of exact position and speed can’t be measured but we can only
give areas of probability in which the electrons can be found.
Area where the
electrons are most
likely to be found
orbiting around the
nucleus
nucleus
Periodic table
Mendeleev – Russian professor of Chemistry 1834 - 1907
Dimitry Mendeleev was a Chemist who created a periodic
table based on elements relative atomic mass and placed
the elements in groups based on the elements similar properties. Not all of
the elements had been discovered at the time he created the table so he left
gaps that has subsequently been filled.
Groups 3 to 12 were
added after
Mendeleev’s table –
these are called the
transition metals
Group 18 – the noble
gases, were not
discovered at that
time and were also
added after.
Periodic Table
The columns (downwards) of a periodic table are called groups.
The rows (across) of a periodic table are called periods.
Elements in the same group all have
the same number of electrons in their
outer (or valence) shells.
Elements in the same period all
have the same number of shells
of electrons in their atoms
37
Group 1 elements
These elements are called the Alkali Metals. They are all very reactive with air and,
especially so, water. The further down the group the more reactive they are. Hydrogen is
not included in this as it does not share similar properties with the rest of the elements.
38
Group 2 elements
These elements are called the Alkali Earth Metals. They all react with air, but are
less reactive than group 1.
39
Group 17 elements
These elements are called the Halogens. They are highly reactive, with
reactivity decreasing down the table. They all contain 7 valence electrons
and readily accept an electron from other atoms.
40
Group 18 elements
These elements are called the Inert Gases (sometimes called Noble gases). They
all have full valence shells and are very unreactive. This property makes them
useful in many practical applications.
41
ions
Atoms with filled outer energy levels are the most stable. An atom will gain
or lose electrons in order to have a filled valence shell.
Ions have different chemical properties and a physical appearance to the
elements they originated from
Cation Sodium (Na)
11+
Sodium now becomes the
sodium ion Na+
Anion Chlorine (Cl)
17+
Chlorine now becomes the
chlorine ion Cl-
ions
Ions are charged particles.
Ions form when atoms gain or lose electrons.
An ion is an atom or
group of atoms which
has gained or lost
electrons.
Elements are most
stable when the outer
energy level (valence
shell) is full.
Elements can lose or
gain electrons when
they react with other
chemicals to form ions
and achieve stability.
ions
Atoms that lose electrons form positively charged ions, or cations.
Atoms that gain electrons form negatively charged ions, or anions.
Cation (Cat)
Anion (an Iron)
+
Metals lose electrons to form
Cations. They have 1-3 electrons in
their outside shell
Non-Metals gain electrons to
form Anions. They have 7-8
electrons in their outside shell.
Ion Chart - Cations
ions
1+
2+
sodium
Na+
magnesium
potassium
K+
iron (II)
ferrous
silver
ammonium
Ag+
NH4+
copper (II)
cupric
Mg2+
Fe2+
Cu2+
zinc
Zn2+
barium
Ba2+
Copper (I)
cuprous
Cu+
Hydrogen
H+
lead
Pb2+
Lithium
Li+
tin
Sn2+
3+
aluminium
iron (III)
ferric
Chromium
Al3+
Fe3+
Cr3+
Ion Chart - Anions
1-
2-
chloride
Cl-
carbonate
CO32-
permanganate
MnO4-
oxide
O2-
thiocyanate
SCN-
sulfide
S2-
iodide
I-
sulfate
SO42-
hydroxide
OH-
sulfite
SO32-
hydrogen carbonate
HCO3-
thiosulfate
S2O32-
hydrogen sulfide
HSO3-
chromate
CrO42-
fluoride
F-
dichromate
Cr2O72-
bromide
Br-
nitrate
NO3-
hypochlorite
OCl-
3phosphate
46
PO4-3
electron configurations of ions – Cations (metals)
The Ca atom has 20
protons and 20 electrons
so has no charge. It is
neutral.
The Ca2+ ion has 20
protons and 18 electrons
so has a 2+ charge.
electron configurations of ions – Anions (non-metals)
The Cl atom has 17
protons and 17 electrons
so has no charge. It is
neutral.
The Cl- ion has 17
protons and 18 electrons
so has a 1- charge.
Compounds
Compounds form from
two or more different
elements bonded
together.
Compounds
The compounds are often more stable than the elements they originated
from and may release this extra energy in the form of heat and/or light when
bonding together.
There are two main types of bonding holding atoms together in a compound;
Ionic and Covalent.
50
Ionic Bonding
Ionic Bonding is where one atom completely takes valence electrons from
another to form ions and the resulting negative and positive ions hold
together with electrostatic attraction. This type of bonding occurs when a
metal and non-metal react and there is a transfer of electrons to form ions.
The ions then combine in a set ratio to form a neutral compound with
negative and positive charges balanced out.
Ionic compounds are the product of chemical reactions between metal and
non-metal ions
Some compounds are ionic compounds, since they are made up of cations
and anions.
Compounds are neutral substances. For ionic
compounds, the charges of the positive ions are
balanced by the charges of the negative ions.
The Anion (F) takes the
electrons off the Cation
(Li) so their outer energy
levels have a stable 8
electrons each.
Anions and Cations have a
strong electrostatic
attraction for each other
so they bond together as a
compound.
Covalent Bonding
Covalent Bonding is where electrons are shared between neighbouring
atoms. This often occurs when two or more non-metals react. No ions are
formed and there is no transfer of electrons. The compound formed is neutral
with no charge.
The valance electrons
(electrons in outside
energy level) are involved
in bonding. These
electrons orbit in pairs.
The negative charge of
the electron pair will
attract the positively
positive nucleus of other
atoms, and this holds the
atoms together in a
molecule.
Covalent Bonding
Extra
for
experts
All covalent bonds are strong. That is it requires a large amount of energy to ‘break’
the bond. However, some covalent bonds are stronger than others. The greater the
overlap of valence orbitals (the area the valence electrons orbit the nucleus) the
stronger the bond.
Covalent Bonding
The electron-pair must lie between the nuclei for the attraction to outweigh
the repulsion of the two nuclei. This ‘sharing’ of electrons between atoms
creates a covalent bond – giving both atoms the stability of a full outer shell.
Covalent bonds are normally formed between pairs of non-metallic atoms.
Some covalent bonds involve
only one pair of electrons and
are known as single bonds.
Other covalent bonds can
involve two pairs of electrons;
double bonds and three pairs of
electrons; triple bonds.
Naming compounds
Lavoisier – French Chemist 1789
Lavoisier devised a system of naming compounds based on their chemical
composition. If the compound is formed between a Metal cation (+ve) and a
Non-Metal anion (-ve), then the compound name joins the two names
together with the metal name first. Names of the ions need to be
remembered.
Sodium
+
hydroxide
Chemical compound formula
1. Write down the ions (with charges) that react to form the compound.
Cation comes before Anion.
Al3+
O2-
2. Cross and drop the charge numbers.
3. Place brackets around a compound ion.
Al2O3
4. If the numbers are both the same remove.
5. If any of the numbers are a 1 they are removed
6. Remove any brackets if not followed by a number
H+
SO4-2
H2(SO4)1
H2SO4
Drawing chemical compounds
G Lewis – American Chemist 1916
G Lewis devised a system of drawing covalent molecules showing
arrangement of atoms and valence electrons – both those involved in
bonding and those that are not (called lone pairs). Electrons in inner shells are
not involved in bonding. These diagrams are called Lewis diagrams. The Lewis
diagram is drawn so that each atom has eight electrons associated with it
(except for hydrogen which has two). This is the octet rule.
Lewis diagram of H2O (water)
Hydrogen
electron
Bonded pair
x
H O
Oxygen electron
x
H
Lone pair
Lewis Diagrams
1. Calculate valence electrons of all atoms. If the molecule is an ion then
subtract the charge from the total electrons and place the charge outside
of square brackets of the Lewis diagram at the end. Example carbon
dioxide.
C=4
CO2
O=6
O=6
16
2. Write down number of pairs of electrons.
16 / 2 = 8 pairs
3. Place atom with least filled valence shell in the centre with the other
atoms arranged around the out side (periphery)
O C O
Lewis Diagrams
4. Bond all atoms together (either x or
O
C
O
= one pair of electrons)
8 pairs – 2 pairs =
6 pairs remaining
5. Place remaining e- pairs around the periphery atoms so each has 4 pairs (including
bond pair) around it.
xx
x
xO
xx
xx
C
6 pairs – 6 pairs =
0 pairs remaining
O xx
xx
6. If there any remaining pairs place them around the outside of the central atom.
7. Rearrange lone pairs (pairs not bonded) into bonded pairs if the central atom does
not have 4 pairs around it.
Rule of orbitals – exceptions to the rule
If there are extra Lone Pairs of electrons left after all of the periphery atoms are filled
in accordance with the octet rule then they are placed around the central atom(s)
according to the Rule of Orbitals.
The Rule of Orbitals: the total number of lone pairs and bond pairs (LP+BP) associated
with an atom cannot exceed the number of Valence Shell Orbitals (VSO = n2,
where n is the row of the Periodic Table in which that atom resides).
n = 1 (H): maximum VSE pairs (LP+BP) = VSO = 1;
n = 2 (B, C, N, O, F): maximum VSE pairs (LP+BP) = VSO = 4 ("octet rule")
n = 3 ((Al, Si, P, S, Cl): maximum VSE pairs (LP+BP) = VSO = 9 etc.
Boron and
Beryllium often
are found with
only 3 lone +
bonded pairs
around them
F
B
F
F
Lewis Diagrams






Extra
for
experts
The number of covalent bonds an atom forms is called its valence.
Some atoms have fixed valence. E.g: H = 1, C = 4, F = 1. (most halogens =
1)
Some atoms have variable valence. For example:
O = 2 (sometimes 3), B, N = 3 (sometimes 4).
an atom bonded to only one other atom is peripheral (monovalent
atoms such as H and F are always peripheral).
an atom bonded to two or more other atoms is central.
Often, the formula is written to indicate connectivity. For example: HCN =
H bonded to C, C bonded to N, H and N are not bonded.
Predicting molecular shapes
Sidgwick and Powell – 1940
Sidgewick and Powell devised a theory to predict the shapes molecule
formed. It is based on the following ideas:
Each electron pair is a region of negative charge
Negative charges repel each other
Electron pairs will be spaced as far apart as possible around a central atom.
This theory is called the Valance Electron Pair repulsion theory.
Three pairs
Four pairs
Predicting molecular shapes
The shapes of molecules are determined by the way the regions of negative
charge are arranged around the central atom in the molecule. A region may
consist of one lone pair of electrons or one bonded pair or two bonded pairs
or three bonded pairs. All of these electron arrangements occupy the same
region of space
Shapes – two clouds
Since regions of electrons are negatively-charged, they repel each other as far apart as
possible. Two clouds arrange themselves on opposite sides of the central atom.
The bond angle will be 180°.
The shape name is linear.
Bonded pair
of electrons
(one cloud)
atom
180°
65
Bond angle
Shapes – three clouds (0 lone pairs)
Three regions of negative charge will cause a bond angle of 120° as they repel each
other.
All the atoms still lie on a flat plane (like a sheet of paper).
The shape is trigonal planar. (or triangular planar)
Periphery atom
120°
Central atom
Shapes – two clouds (1 lone pair)
When one of the clouds of electrons is a lone pair it will have a slightly greater push to
the bonded pairs. This is because the lone pair are only orbiting around one positive
nucleus and their negative charge is less ‘neutralised’ than if they had another nucleus
to orbit around. The regions of negative charge repel to a trigonal planar shape. The
bond angle between the remaining pairs is approximately 120° .
The final shape formed by the atoms is called bent.
Lone pair
Bonded pair
~120°
Shapes – four clouds (0 lone pairs)
When four regions of negative charge are around a central atom they repel each other
into a 3-dimensional shape. The bond angle is now 109.5°. This is because it is a
sphere divided into 4 rather than a circle.
This shape is tetrahedral.
109.5°
Shapes – four clouds (1 lone pair)
The four regions of negative charge still occupy a 3-dimensional tetrahedral shape.
(The lone pair, however, exerts a stronger repulsion to the remaining bonded pairs).
The bond angle is 109.5°.
The final shape the bonded atoms form is a trigonal pyramid (or a triangular pyramid)
Lone pair
109.5°
Shapes – four clouds (2 lone pairs)
The 4 regions of negative charge repel each other to a (warped) tetrahedral shape. But
The two lone pairs create a much stronger repulsion than one lone pair and the bond
angle between the remaining bonded pairs is smaller again at approximately 109.5°
(compared to 120 ° of the bent shape with only 3 regions of negative charge). The final
shape the bonded atoms form is called Bent.
Lone pair
109.5°
Shape Summary
Electron clouds (lone
pairs or bond groups)
2
Linear
CO2
180°
3
4
no lone pairs
1 lone pair
no lone pair 1 lone pair 2 lone pair
triangular planar
Bent
Tetrahedral
triangular
pyramid
CH4
109.5°
NH3
109.5°
BF3
120°
SO2
~120°
bent
H2O
109.5°
Drawing Shapes
1.
Draw molecule – Lewis
diagram first
2.
Calculate number ofregions
of negative charge around
central atom
•
•
Single, double or triple
electron bonds occupy
1 cloud
Note: make
sure the
question asks
you to draw
a shape and
not a lewis
diagram
receding
Same plane
A lone pair of electrons
occupy 1 cloud
3.
Calculate how many atoms
are joined to the central
atom
4.
Name / draw shape
approaching
Discussing shapes questions – NCEA example
Explain why the shape of the CO2 molecule is linear but the shape of H2O is
bent?
1. The C (central atom) of CO2 has 2 regions of negative charge around it in
the form of double bonds connected to a O atom. (draw lewis diagram)
2. Each of the regions of negative charge repel each other the furthest away
from each other in 3 dimensional space into a linear.
3. There are no lone pairs so the final CO2 molecule therefore also forms a
linear shape
1. The O molecule (central atom) of H2O has 4 regions of negative charge
around it in the form of two single bonds connected to a H atom and two lone
pairs. (draw lewis diagram)
2. Each of the regions of negative charge repel each other the furthest away
from each other in 3 dimensional space and form a tetrahedral shape.
3. However with only 2 of the regions bonded to atoms the final shape the
H2O molecule forms is a bent shape
Electronegativity
Electronegativity is the attraction that an atom has towards electrons from
another atom. The greater the electronegativity the stronger the pull it has
towards other electrons.
Trends in the periodic table
 The larger the nucleus (with the positive protons) the stronger the
electronegativity, this means it increases from left to right.
The further the valence electrons are from the nucleus the lesser the
electronegativity, therefore the electronegativity decreases down a group.
Electronegativity
We use a pauling scale to determine electronegativity. The scale starts close
to 0 – with minimal electronegativity and goes up to 4 with the highest
electronegativity. Most of the Inert gases do not have a value because of
there non reactivity with other atoms.
Extra
for
experts
Ionic – covalent bond continuum
Bond types between atoms can depend on the electronegativity of
the atoms. Rather than discrete categories, molecules fall along a
continuum
Covalent
0—0.4
Polar Covalent
0.4—1.6
Ionic
>1.6
Polarity
If two identical atoms are bonded together then they have exactly the same
amount of attraction to the shared electrons in the bonded pair. This is
because their electronegativity is the same. This becomes a non-polar
molecule with non-polar bonds. Example - Iodine molecule I2
Path of electrons
Evenly shared
If two different types of atoms are bonded together then they will exert
different levels of attraction for the orbiting electrons. That is because they
may have different numbers of electron shells and different numbers of
protons in their nucleus. This will cause an electronegativity difference. These
bonds become polar bonds. Example – hydrochloric acid HCl
Slightly positive
Electrons orbit less
δ+
δ-
Slightly negative
Electrons orbit more
Polarity
δ-
δ+
Cl
H
Polarity may also be shown as an arrow, with a cross, +ve, at the tail. The
arrow head is the –ve end.
Polarity
If two bonded atoms are the same, the
bond is said to be non-polar. i.e. I2
The whole molecule is also non-polar
because there is no electronegativity
difference and the valence electrons orbit
each atom evenly. If two different atoms
are bonded they form a polar bond, as
there is an electronegativity difference and
the valence electrons spend more time
around the atom with the higher
electronegativity value (that atom becomes
slightly negative ) The atom that the
valence electrons spend less time around
becomes slightly positive.
Symmetry and Polarity
The polarity of a molecule with
polar bonds depends upon
whether the molecule is
symmetrical or not.
A symmetrical molecule (one
where the centres of peripheral
atoms coincide) becomes a
non-polar molecule – as the
charges balance out
An unsymmetrical molecule
(one where the centre of
peripheral atoms do not
coincide) is a polar molecule.
Answering NCEA Polarity Questions
Explain why molecules x (CCl4) and y (NCl3) are polar and non-polar?
Polar molecule
1. molecule (NCl3) is polar (state which
one)
2. (NCl3) contains polar bonds due to
electronegativity difference of N and Cl.
3. over the whole molecule the atoms
are not distributed symmetrically in 3
dimensions because its shape is (state
which one) and has lone pairs of
electrons
4. polar bonds do not cancel each other
out and the whole molecule is polar.
Non-polar molecule
1. molecule (CCl4) is non-polar (state
which one)
2. (CCl4) contains polar bonds due to
electronegativity difference of C and Cl.
Cl attracts more electrons than C
because it has a bigger atomic number
than C but with the same number of
shells
3. over the whole molecule the atoms
are distributed symmetrically in 3
dimensions because its shape is (state
which one)
4. polar bonds cancel each other out and
the whole molecule is non-polar.
Solubility
The solubility of a substance is the amount of that substance that will dissolve in a
given amount of solvent. Solubility is a quantitative term. Solubility's vary
depending on the solvent and the solute. The terms soluble and insoluble are
relative. Some substances can be sparingly soluble where only the most minute
percentage dissolves. For a solute to dissolve the attraction to the solvent
molecules must be stronger than the bonds holding the atoms/molecules of the
solute together.
GZ Science Resources 2013
Water as a solvent
Solute
(salt)
Solvent
(water)
Solution
(saltwater)
A solution is made up of a
solvent and a solute. A solvent is
a substance such as water that is
able to dissolve a solute. The
solvent ‘pulls apart’ the bonds
that hold the solute together and
the solute particles diffuse
(spread randomly by hitting into
each other) throughout the
solvent to create a solution. The
solution is a mixture with evenly
spread solvent and solute
particles. These particles can be
physically separated by
evaporation.
Solutions form when a solute is dissolved in a solvent
When a solid mixes into a liquid and can longer be seen it has dissolved. The liquid is
called the solvent and it pulls apart the bonds between the solid particles, called the
solute, and they diffuse. A solution is then created when the solvent particles (often
water) are mixed up with the broken apart solute particles.
Solubility
For a solute to dissolve, the solvent particles must
form bonds with the solute particles that are of
similar strength, to the bonds between the solute
particles.
Water, being polar attracts ions because they are
charged and so dissolves many ionic substances.
Polar Solvents
The water molecule has two
polar bonds. Due to the
asymmetry of the molecule,
their polarities reinforce making
the oxygen side of the molecule
partially negative (δ-) and the
hydrogen side partially positive
(δ+). Such molecules are called
‘polar’. The separation of
charge in the molecule (δ+ - δ) is called a ‘dipole’.
Polarity causes a stream of
water molecules to attract to a
charged plastic pen.
Dissolving and Polarity
Polar substances dissolve polar substances.
e.g. Water, being polar attracts the molecules of
other polar substances (e.g. HCl) and will dissolve
them.
Polar substances will not dissolve non-polar
substances.
e.g. Water, (polar) has a stronger attraction to
itself than to non-polar molecules (e.g.
cyclohexane) and will not dissolve them.
Non-polar substances dissolve non-polar
substances.
e.g. Non polar solvents (like cyclohexane) attract
non-polar solutes (like napthalene) by the same
weak Van der Waals forces they attract
themselves by and so will dissolve non-polar
solutes.
Ionic solid dissolving in water
Common Polar and Non-polar molecular substances
Polar
Non-Polar
water
methanol
ethanol
acetic acid
hydrogen
chloride
cyclohexane
benzene
hydrocarbons (e.g.
petrol)
oxygen
hydrogen
nitrogen
iodine
Sample NCEA Questions:
Potassium chloride will not dissolve in non-polar solvents, but will dissolve
in water. Explain by relating the property to the structure and bonding
within the solid.
Use the structure and bonding in H2O and SO2 to explain why SO2 is soluble
in H2O.
Groups of substances
Substances are grouped together according to the type of bonds they have between
particles.
This year will cover four groups of substances; Molecular, metallic, ionic and covalent
network. The physical properties of these groups will be linked to their structure.
Molecular solids
Ionic solids
Non-metals forming molecules
Non-metals and metals together
forming a ionic compound
S2
sulfur
HCl
Hydrogen chloride
I2
iodine
KI
Potassium iodide
CuSO4
NaCl
Copper sulfate
Sodium chloride
Metallic solids
Covalent network solids
Elements that are metals
Carbon and silicon dioxide
Fe
iron
SiO2
Silicon dioxide
Al
aluminium
Cu
copper
C
diamond
C
graphite
Non-polar Molecular solids
non-metal + non-metal
Molecules are held together by
weak intermolecular forces
caused by temporary dipoles
induced by electrons randomly
spending more time around one
nucleus than the other.
Within the Molecules, the atoms
are held together by strong
covalent bonds.
Weak intermolecular bond
Strong covalent bond
Polar Molecular solids
Polar molecules held together by weak inter molecular forces caused by
permanent dipoles induced by electrons spending more time around one
nucleus in the molecule that has greater electronegativity than the other. The
δ –ve end of one molecule is attracted to the δ +ve end of another.
δ +ve
δ -ve
δ +ve
δ +ve
δ +ve
δ -ve
δ +ve
91
δ -ve
δ +ve
Weak intermolecular forces
Weak intermolecular forces of attraction
3 kinds
instantaneous dipole – induced dipole in non-polar molecules
permanent dipole – permanent dipole in polar molecules
hydrogen bonding – strong dipoles for example in water between H and O
Note the distinction:
Intra-molecular Forces: the strong
bonding forces within a molecule. i.e. the
covalent bonds holding the molecule
together.
Inter-molecular Forces: the weak
bonding forces between molecules due
to the attractions between partial
charges. i.e. permanent dipole
Polar Molecular solids - solubility
1. Hydrogen chloride (HCl)
is a molecular solid
2. Hydrogen chloride is
made up of covalently
bonded atoms to form
molecules
3. these molecules are held
together by weak
intermolecular forces
4. these molecules are
polar therefore the
electrostatic attractions of
water molecules (which is
stronger than the weak
intermolecular forces) have
sufficient strength to pull
the molecules apart hence
hydrogen is soluble
Polar molecules
H
Cl
Cl
H
H
Cl
δ -ve
H
δ +ve
O
H
δ +ve
Non-polar Molecular solids - solubility
1. Iodine is a molecular
solid
2. Iodine is made up of
covalently bonded atoms
to form molecules
3. these molecules are held
together by weak
intermolecular forces
4. Iodine is non-polar
therefore the electrostatic
charges of the water do
not have sufficient strength
to overcome the weak
intermolecular forces
holding the molecules
together hence iodine is
insoluble
+
+
δ -ve
H
δ +ve
O
+
+
H
δ -ve
O
δ +ve
H
δ +ve
Non- Polar molecules
H
δ +ve
Molecular solids – Melting point
For example:
1. Carbon dioxide is a molecular solid (at low temperatures below -56◦C)
2. Carbon dioxide is made up of covalently bonded atoms to form molecules
3.these molecules are held together by weak intermolecular forces
4. these forces require small amounts of energy to break apart the solid (but not the
individual molecules which are held together by strong covalent bonds) therefore carbon
dioxide has low melting point
O
Heat
O
C
O
C
O
Many molecular
solids are only
solid at
temperatures
well below 0◦C
and at room
temperature they
are gases
Molecular solids - hardness
For example:
1. sulfur is a molecular solid
2. sulfur is made up of covalently bonded atoms to form molecules
3. these molecules are held together by weak inter molecular forces
4. these forces require small amounts of energy to break apart the solid (but
not the individual molecules which are held together by strong covalent
bonds) therefore sulfur is easily broken up
S
S
S
S
S
S
Force
applied
S
S
Weak intermolecular force
Molecular solids - Conductivity
In order for a substance to be electrically conductive there must be free moving
charged particles
For example:
1. Iodine is a molecular solid
2. Iodine is made up of covalently bonded atoms to form molecules
these molecules are held together by weak intermolecular forces
3. there are no free moving charges therefore iodine cannot conduct electricity
nucleus
+
+
Valence
electrons
+
Covalently
shared
electrons
Weak intermolecular bonding
+
Fully occupied valence
electrons remain in
‘fixed orbit’ around
nucleus and are not
available to carry
charge. The molecule
is neutral
Metallic Solids - structure
Metals atoms are arranged as positive ions held in place in ordered layers by
strong attractive non-directional bonding, forming a lattice. - this gives metals
strength.
Metal atoms are held together in a
3–D lattice by metallic bonding in which
valence electrons are attracted to the
nuclei of neighbouring atoms. The
attraction of the metal atoms for the
valence electrons is not in any particular
direction; therefore metal atoms can move
past one another without disrupting the
metallic bonding, therefore metal is
ductile.
The atoms are packed tightly together - this makes metals dense
Metallic Solids - Conductivity
Electrons from the outer shells of the metal atoms move freely throughout the
lattice. - this makes metals excellent conductors of heat and electricity
99
Metallic Solids - Conductivity
Free moving charged particles are required to carry a charge and for a substance to be
electrically conductive
For example:
Always state what type of solid a substance is
first.
1. copper is a metallic solid
2. copper is arranged as positive ions held in place in ordered layers by strong
attractive non-directional forces, in a sea of de-localised electrons
3. electrons are free moving hence can carry a charge
4. therefore copper can conduct electricity
Metallic Solids - Solubility
For example:
1. lead is a metallic solid
2. Lead is arranged as positive
ions held in place in ordered
layers by strong attractive nondirectional forces, in a sea of
de-localised electrons
3. these forces require a large
amount of energy to break
therefore the electrostatic
attractions of water molecules
do not have sufficient strength
to pull the atoms apart
4. therefore lead is insoluble
In order for substance to dissolve in water (a
polar liquid) the attraction between the
particles in a substance must be less than the
attraction towards water molecules
Metallic Solids – Malleability and ductility
For example:
1. iron is a metallic solid
2. iron is arranged as positive ions held in
place in ordered layers - a lattice, by
strong attractive non-directional forces, in
a sea of de-localised electrons
3. these forces require large amounts of
energy to break apart the solid therefore
aluminium is not easily broken up
4. However Layers can slide over each
other, and as the attractive forces are nondirection the metallic particles remain
strongly bonded. – this gives the metallic
solids the properties of being malleable
(moulded into flat sheets) and ductile
(drawn out to thin wires)
Layers of ions can slide over each
other without breaking- this
makes metals hard and also
malleable and ductile
Metallic Solids – melting point
The strength of the bonds between particles determines the energy required to
break them, and therefore the amount of energy to change a solid into a liquid (the
melting point) where the bonds are somewhat broken.
Metals in general, have very strong bonds which makes them solid at room
temperature (Mercury is the exception)
For example:
1. Aluminium is a metallic solid
2. Aluminium is arranged as positive ions held in place in ordered layers by
strong attractive non-directional forces, in a sea of de-localised electrons
3. These forces require a large amount of energy (high temperature) to break
apart the metallic solid therefore the melting point is very high.
Three steps to answering structure and physical properties questions.
The first is state the name of the solid.
The second is describe the structure of the solid.
The third is link the structure of the solid to the physical property discussed.
Ionic Solids - structure
Metal + Non-Metal
An ionic solid is made up of ions
held together by strong
electrostatic forces (ionic
bonding) between +ve (cations)
and –ve (anions) ions in a 3-d
lattice.
Ionic Solids - Solubility
In order for substance to dissolve in water (a polar liquid) the attraction between the
particles in a substance must be less than the attraction towards water molecules
For example:
1. Sodium chloride (NaCl) is an ionic solid
2. Sodium chloride is made up of ions held together by strong electrostatic attractions
between +ve and –ve ions in a lattice
3. the electrostatic attractions of water molecules have sufficient strength to pull the
ions apart however
4. therefore the solid will dissolve and is soluble
NaCl first place in water
Na+ and Cl- ions breaking apart
The positive hydrogen end of
water is attracted to the
anions and the negative
oxygen end of water is
attracted to the cations
Ionic Solids - Conductivity
Free moving charged particles are required to carry a charge and for a substance to be
electrically conductive
For example:
1. Sodium chloride is an ionic solid
2. Sodium chloride is made up of ions held together by strong electrostatic
forces between +ve and –ve ions in a 3-d lattice
3. when solid the ions are not free to move therefore it doesn’t conduct
electricity
4. However when melted, or dissolved in solution, the bonds are broken and the
ions are free to move therefore sodium chloride can conduct electricity
Na+
Cl-
Na+
Cl-
Cl-
Na+
Cl-
Na+
Distilled water does not
conduct a current
Positive and negative ions fixed in
a solid do not conduct a current
In solution, positive and negative
ions move and conduct a current
Ionic Solids – Hardness and brittleness
For Example:
1. Sodium chloride is an ionic solid
2. Sodium chloride is made up of ions held together by strong electrostatic
attractions between +ve and –ve ions in a 3-d lattice so requires a lot of energy
to break the bonds
3. However if sideways force is applied and a sheet of the lattice slides then ions
of the same charge may come in contact with each other and repel hence the
ionic solid is brittle (and can break into pieces)
Na+
Cl-
Na+
Cl-
Na+
Cl-
Na+
Clforce
Cl-
Na+
Cl-
Na+
Cl-
Na+
Cl-
Na+
Ionic Solids – Melting Point
For example:
1. sodium chloride is an ionic solid
2. Sodium chloride is made up of ions held together by strong electrostatic
attractions between +ve and –ve ions in a 3-d lattice
3. Because these strong bonds require a large amount of energy to break the
ionic solids have a high melting point.
Na+
Cl-
Na+
Cl-
Cl-
Na+
Cl-
Na+
Covalent Network Solids - structure
All atoms are held together by
strong covalent bonds
Diamond is a 3-dimensional
covalent network structure
where atoms are held together
by strong covalent bonds in all
planes
diamond
graphite
Graphite is a covalent network
structure that is in 2 dimensional
sheets (graphite). Between the
layers are free moving electrons
from the valance electrons of the
carbon atoms.
Silicon dioxide (SiO2) is a 3dimensional covalent network
structure
Silicon dioxide
Covalent Network (3D) - Conductivity of diamond
For example:
1. diamond is a 3dimensional
covalent network
structure
(diamond)
2. all atoms are
held together by
strong covalent
bonds
3. there is no free
moving charged
particles
4. therefore
diamond cannot
conduct electricity
Covalent Network (2D) - Conductivity of graphite
For example:
1. graphite is a covalent
network that is in 2
dimensional sheets
2. between the layers
are free moving
electrons from the
valance electrons of the
carbon atoms.
3. the free moving
electrons can carry a
current
4. therefore graphite
can conduct electricity
Covalent Network Solids - Solubility
For example:
1. Silicon Dioxide is a 3dimensional (or 2-dimensional)
covalent network structure
2. all atoms are held together by
strong covalent bonds
3. these forces require a large
amount of energy to break
therefore the electrostatic
attractions of water molecules
do not have sufficient strength
to pull the ions apart
4. hence silicon dioxide will not
dissolve in water and is insoluble
Covalent Network Solids - Melting Point
For example:
1 Diamond is a 3-dimensional (or 2-dimensional) covalent network structure
2. all atoms are held together by strong covalent bonds
3. these forces require a large amount of energy to break
4. therefore diamond has a very high melting point.
Covalent Network Solids - Hardness
For example:
1. Diamond is a 3-dimensional (or 2-dimensional) covalent network structure
2. all atoms are held together by strong covalent bonds
3. these forces require a large amount of energy to break
4. therefore diamond is very hard.
Solids Summary
Name of solid
substance
Type of
particle in
solid
Attractive
force broken
when solid
melts
Attractive
force
between
particle –
weak or
strong
(hardness)
Relative
melting point
solubility
Electrical
conductivity
molecular
molecules
Weak
intermolecular
weak
low
Yes if polar
No if nonpolar
no
metal
atoms
Metallic
bonding
strong
high
no
yes
ionic
ions
Electrostatic
strong
high
yes
Only if
molten or
in solution
strong
high
no
no
Ionic
bonding
covalent
atoms
Covalent
bonding
Property
Type of Solid
Molecular
Solubility in
Water
Electrical
Conductivity
Melting
Point
Hardness
Covalent network Ionic
Metallic
Energy Changes
Enthalpy and Enthalpy Change
∆H
Enthalpy (or Heat Content) is the energy in a substance due to kinetic energy
of particles and potential energy in chemical bonds
Enthalpy change
∆H is the
difference in enthalpy of
products HP and reactants HR
∆H = HP - HR
The unit for Enthalpy is kiloJoules (kJ)
Enthalpy Change
HP (products) and HR (reactants) cannot be measured.
We can measure Enthalpy change ( ∆H ) by measuring energy;
Released to surroundings
(Exothermic Reactions)
Absorbed from surroundings
(Endothermic Reactions)
119
Exothermic Reactions
These are reactions where heat energy
is released into the surroundings.
Surroundings gain heat energy.
(increase in temperature )
Products will have less energy than
reactants.
∆H is NEGATIVE (-)
Endothermic Reactions
These are reactions where heat
energy is absorbed from the
surroundings.
Surroundings lose heat energy.
(Decrease in temperature)
Products will have more energy than
reactants.
∆H is POSITIVE (+)
Exothermic reactions
Any combustion reaction is exothermic. The bonds holding the atoms of fuel
molecules together (usually consisting of carbon and hydrogen atoms)
release a lot of energy in the form of light and heat when they are broken.
The total energy holding the bonds together in the products are less than the
total energy in the reactions and the difference is released.
GZ Science Resources 2013
121
Endothermic reactions
Melting ice is an example of
an endothermic reaction. The
solid ice (water) atoms that
are in a fixed pattern are
barely moving and need to
absorb energy in order to
move faster and break the
bonds to form water in a
liquid state.
GZ Science Resources 2013
122
Enthalpy Diagrams
Endothermic Reactions
e.g. Reacting methane with steam at
high pressure and temp. Energy is
absorbed
Exothermic Reactions
e.g. Burning of methane in air. Energy is
released
Enthalpy Change
An exothermic reaction will release
energy and the products will be at a
lower enthalpy level than the
reactants.
The reaction system will feel will feel
hot to the touch as the energy is
released as heat energy.
An endothermic reaction will absorb
energy and the products will be at a
higher enthalpy than the reactants.
The reaction system will feel cool to
the touch as heat energy is taken from
the surroundings, including your skin,
and used to break bonds in the
molecules.
Energy Diagrams
Endothermic Reaction e.g. Reacting methane with steam at high pressure and
temp. Energy is absorbed
Energy Diagrams
Exothermic Reaction e.g. concentrated Hydrochloric acid reacting with zinc metal
HR
HP
Breaking Bonds - endothermic
Bonds holding atoms and molecules together require the input of energy in order to
break them apart therefore breaking of bonds is an endothermic reaction. The input
of energy (usually light or heat energy) cause the atoms and molecules to move faster
and ‘pull away’ from each other. Each type of bond has its own specific amount of
energy, called bond energy measured in kJ, required to break its bond.
Forming Bonds - exothermic
Bonds forming between atoms and
molecules release energy therefore bond
forming is an exothermic reaction. Bonds
are formed to form a stable molecule.
If more energy is required to break the bonds of the
reactants than released when the bonds of the
products are form then the overall reaction is
endothermic.
If less energy is required to break the bonds than is
released when the bonds of the products are formed
then the overall reaction is exothermic.
Enthalpy in Dissolving
If more energy is released when water bonds to the solute than it takes to
separate the solute, the dissolving is exothermic and the temperature
increases. An example is adding a strong acid (such as sulfuric acid) or base
(such as sodium hydroxide)
Standard conditions
Measurements depend on conditions
When measuring a enthalpy change
you will get different values under
different conditions. For example, the
enthalpy change of a particular
reaction will be different at different
temperatures, different pressures or
different concentrations of reactants.
The values for enthalpy are given for
standard conditions, indicated by the
superscript θ
Standard conditions include:
Temperature of 25°C
Atmospheric pressure conditions of
1ATM
Concentration of 1mol per Litre
Thermochemical Equations
Exothermic
CH4 + 2O2
Endothermic
CO2 + 2H2O
CH4 + H2O
∆H = -888kJmol-1
CO + 3H2
∆H = 206 kJmol-1
This thermochemical equation reads;
888kJ of heat is released when 1 mole of
CH4 reacts with 2 moles of O2 to produce
1 mole of CO2 and 2 moles of H2O
Use thermochemical equations to find
This thermochemical reaction reads;
206kJ of heat is absorbed when 1 mole
of CH4 reacts with 1 mole of H2O to
produce 1 mole of CO and 3 moles of
H2.
∆H, n and m.
n = m/M
n = moles (6.02 x 1023 particles) m = mass (grams) M = Molar Mass (gmol-1 )
Thermochemical Equation Example
CH4 + 2O2
CO2 + 2H2O
∆rH = -888kJmol-1
Use the equation above to find heat released if 2.5 moles of CH4 burns.
1 mole of CH4 releases 888kJ
2.5 moles CH4 releases x kJ
x = 2.5 x 888 = 2220kJ
An equation is a mole ratio – the number in
front of each substance tells you how many
moles of that there is to any other
substance.
For example there is 1 mole of CH4 to every
2 moles of O2
The enthalpy of the equation shows you the
amount of energy per unit of substance.
888 = 1CH4 888 = 2O2 (444 = 1O2)
Thermochemical Equation Example 2
CH4 + 2O2
CO2 + 2H2O
∆rH = -888kJmol-1
Calculate the amount (in moles) of H2O produced when the reaction above
releases 10,000kJ.
An alternative method is to find
out how much energy is released
per mole first
2 moles H2O = 888kJ
Therefore 1 mole H2O = 444kJ
Amount of
mols in
equation
10,000kJ/444kJ = 22.5
Amount energy per
unit of substance
Total energy
released
So 22.5 moles of water are
produced at 444kJ to reach
10,000kJ
(22.5 x 444 = 10,000)
Thermochemical Equation Example 3
CH4 + H2O
CO + 3H2
∆rH = 206kJmol-1
M(C) = 12gmol-1 M(O) = 16gmol-1
Calculate the energy required to produce 1kg of CO gas from the reaction
above
Step one
moles of CO produced
M = 1000g M(CO) = 28gmol-1
n = m/M
n = 1000/28 = 35.7 moles
1kg = 1000g. Must be converted to
grams first
If Molar mass is not given then use the
periodic table
The units are kJ not kJmol-1 as it is total
amount not amount per mole.
Step two
1 mole CO produced requires 206kJ (as per the equation above)
35.7 mols CO produced so…..
enthalpy = 35.7 x 206 = 7354kJ
Bond Enthalpy
The high values for bond enthalpy explains why some substances are very resistant to
chemical attack and form very stable molecules
In a polyatomic (more than one atom) molecule, the bond strength between a given
pair of atoms can vary slightly from one compound to another. The value given for
bond enthalpy is the average of all these variations.
A multiple bond (double/triple) is always stronger than a single bond because more
electrons bind the multiple bonded atoms together.
The table below shows some common average bond enthalpies.
Bond enthalpy /kJ mol-1
H-H
436
Bond enthalpy /kJ mol-1
C-H
Bond enthalpy /kJ mol-1
412
C=C
612
837
H-O
463
C - Cl
338
C=oCC
C
H-N
388
C-F
484
C=O
743
H - Cl
431
C-O
360
O=O
496
H-F
565
C-C
348
NN
944
F-F
158
O–O
146
Cl - Cl
242
Bond Energies
Bonds Broken – Endothermic
Bonds formed – Exothermic
∆rH°= ∑ (energy of bonds broken) - ∑(energy of bonds formed)
Note: Bond energies calculated for gases. Convert using ∆vapH° or ∆subH° if in solid
or liquid state.
CO(g) + H2O(g) → H2(g) + CO2(g)
Bonds Broken
CΞO
H-O x 2
995kJ
2(463)kJ
1921kJ
∆rH°= -1.0 Kjmol-1
Bonds formed
C=O x 2
H-H
2(743)kJ
436kJ
1922kJ
∆rH°= 1921kJmol-1 – 1922kJmol-1
The equation can also be arranged
∆rH°= -1.0 kJmol-1
to calculate unknown bond energy
Bond Energies
Bonds Broken – Endothermic
Bonds formed – Exothermic
∆rH°= ∑ (energy of bonds broken) - ∑(energy of bonds formed)
Reactants:
Draw lewis
diagrams to
calculate the
number and
type of bond
Multiply the bond energy given by the
number of bonds
Products:
Draw lewis
diagrams to
calculate the
number and
type of bond
Total the
bond energy
for product
molecules
Total the
bond energy
for reactant
molecules
bonds broken (reactants) minus bonds
formed (product) = total enthalpy
Using Bond Enthalpies to calculate ∆rH°
Bond enthalpy is the change in enthalpy when the covalent bond, in a gaseous
molecule, is broken. It is always a positive value because bond breaking always
requires an input of energy.
Making bonds releases energy so generally speaking the more bonds a substance
can form the more stable it will be.
The strength of a covalent bond depends on the electrostatic attraction between
the positive nuclei and the shared electron pair. The larger the atomic radius of
an atom (which increases down a group) the further the shared electron pair
from the positive nucleus – which creates decreasing electrostatic attraction.
Therefore the weaker the covalent bond and the lower the value of bond
enthalpy
The stronger a covalent bond, the higher the value of the bond enthalpy. The
units are kJ mol-1.
Calculating ∆rH° given the standard heats of formation of reactants and
products.
The standard enthalpy of any reaction can be obtained by subtraction of the standard
enthalpies of formation of reactants from those of the products.
rHo =  n fHoproducts -  n fHoreactantss
where n is the stoichiometric coefficient of each substance in the reaction equation.
Example
Using the standard heats of formation of CO2(g), H2O(l), and C6H12O6(s), calculate the
standard enthalpy of combustion of glucose.
fHo(C6H12O6, s) = -1268 kJ mol-1
fHo(CO2, g)
= -394 kJ mol-1
fHo(H2O, l)
= -286 kJ mol-1
fHo(O2, g)
= 0 kJ mol-1
Note - Start by writing an equation for the combustion of 1 mole of glucose.
C6H12O6(s) + 6O2(g)
 6CO2(g) + 6H2O(l)
rHo
=  nfHoproducts -  nfHoreactants
rHo = ( 6 x -394 + 6 x -286) - (1 x -1268 + 6 x 0) = - 2812 kJ mol-1
Calculating enthalpy using temperature change
Extra
for
experts
ΔH can be calculated if you can measure the temperature change (°C) of a
particular mass (g)
ΔH = m x c x ΔT
ΔH = mass (g) x ΔT (°C) x specific heat capacity (c) (c = 4.18 J g-1 °C)
1. Calculate the mass of reactants – record in grams (one ml liquid = one gram)
2. Measure the temperature of the reactants and the temperature of the products and
calculate ΔT (°C)
3. Calculate ΔH using formula above
Enthalpy of Reaction using calorimetry
Extra
for
experts
To measure enthalpy changes, the reaction it is carried out in an insulated container
(such as a polystyrene cup) and the temperature change (in °C) is measured. Using
this temperature change, ΔT, and the value of the specific heat capacity, c, the
amount of energy transferred to the mass m of substance (usually water) can be
calculated using the expression
∆H = m c ΔT
The specific heat capacity of the water is 4.18 J °C-1 g-1.
Every 1mL of water can be taken as 1g due to its density
1. Calculate the mass of reactants – record in grams
(one ml liquid = one gram)
2. Measure the temperature of the reactants and
the temperature of the products and calculate ΔT
(°C)
3. Calculate ΔH using formula above
Specific Heat Capacity (c)
The heat required to produce a 1oC rise in 1 kg of a substance.
Extra
for
experts
e.g. c(H2O) = 4.18 kJ oC-1kg-1
The energy change when a body of mass m experiences a temperature rise of ΔT is given
by:
Q
= m c ΔT
energy change = mass x specific heat capacity x change in temperature
Example: Calculate the energy change when a 9
kg mass of water increases its
temperature by 20 oC.
Q
= m c ΔT
= 9 kg x 4.18 kJ oC-1kg-1 x 20 oC
= 752 kJ
Note: Watch units g or kg, J or kJ and be ready to
convert if required.
142
Enthalpy Changes
Exothermic - energy released
∆ H negative
Endothermic - energy absorbed ∆ H positive
Enthalpy of reaction (∆rH) for any given reaction
Standard conditions (∆rH°)
1 atmosphere of pressure, 25°C
Enthalpy of fusion (∆fusH°)
1 mol solid to liquid state
Enthalpy of vapourisation (∆vapH°)
1 mol liquid to gas state
Enthalpy of sublimination (∆subH°)
1 mol solid to gas state
solid
∆subH°
∆fusH°
gas
liquid
∆vapH°