CHEMICAL BONDING Part 1 Introduction
to Chemical Bond Formation
Brown's Chemistry Chemical Bonding GCSE/IGCSE/O/AS/A Level Revision Notes
The five linked pages introduce to the concept of
a chemical bond and why atoms bond together, types of chemical bonds and which
electron arrangements are particularly stable leading to stable chemical bonds.
Through the use of dot and cross electronic diagrams is described and there are detailed
notes on ionic bonding i.e. the mutual attraction of oppositely charged ions to
give ionic bonds and the properties of ionic compounds, covalent bonds and the
formation of small simple molecules and their properties, macromolecules like
polymers and giant covalent structures like diamond, graphite and silica.
Finally metallic bonding is described to explain the structure and physical
properties of metals. These notes on chemical bonding are designed to meet the
highest standards of knowledge and understanding required for students/pupils
doing GCSE chemistry, IGCSE chemistry, O
Level chemistry, KS4 science courses and a basic primer for AS/A Level chemistry
courses. These revision notes on the periodic table should prove useful for the
new AQA, Edexcel and OCR GCSE (9–1) chemistry science courses.
Introduction – why do atoms bond together? (this page,
and sub–index for Parts 2–5 (this page)
Ionic Bonding – compounds and properties
Covalent Bonding – small simple molecules and their properties
Covalent Bonding – macromolecules and giant covalent structures
Metallic Bonding – structure and properties of metals
Extra notes on chemical bonding for
ADVANCED A Level Students ONLY (IB, US grade 11-12)
Electronegativity, bond polarity, type of chemical
6.2 More on ionic
structures and ionic bonding
electron configurations for atoms and ions)
6.3 More on
covalent bonding - dative covalent bonding
6.4 Types of
Crystal Structure and their relative physical properties
Shapes of molecules
VSEPR theory - lots of inorganic
Some other molecules/ions of carbon, nitrogen, sulphur
shapes and bond
angles of organic molecules
Intermolecular forces - intermolecular
Introduction to intermolecular forces -
Detailed comparative discussion of boiling points of 8 organic molecules
Boiling point plots for six
homologous series and explaining the trends and differences
Other case studies of
boiling points related to intermolecular forces for a variety of compounds
Evidence and theory
for hydrogen bonding in simple covalent hydrides and its importance in other
NANOSCIENCE – NANOTECHNOLOGY –
NANOCHEMISTRY (index of pages and keyword index)
SCIENCE (alphabetical index at top of page)
Keywords/phrases/names sub–index for Parts 2–5: Examples of ionic compounds described:
sodium chloride NaCl (exemplar for any Li/Na/K + F/Cl/Br/I combination),
magnesium chloride MgCl2 (exemplar for any Mg/Ca +
F/Cl/Br combination), aluminium fluoride AlF3,
potassium oxide K2O (exemplar for any Li/Na/K + O/S
oxide MgO/CaO and magnesium/calcium sulphide (MgS/CaS), aluminium oxide Al2O3
(exemplar for Al2S3) * Examples of covalent molecules: simple small molecule bonding e.g. water *
physical properties of small
molecules * giant network bonding
– giant molecules e.g. carbon C–diamond/graphite, silicon Si/silica SiO2 *
properties of giant covalent structures *
* properties of polymers *
hydrogen chloride HCl,
methane CH4, oxygen
O2, carbon dioxide CO2,
N2, ethane C2H6,
methanol CH3OH, carbon
(diamond), carbon (graphite), carbon
(buckminsterfullerene/fullerenes), silica/silicon dioxide
SiO2 * examples of ionic compounds
properties of ionic compounds * If your ionic compound is not listed, look for a compound with a similar
formula and you should be able to work it out from the example given. The use
of the word exemplar implies you are dealing with the same set of outer electron
arrangements (configurations), which is why you can work out lots more dot
and cross diagrams of ionic compounds by understanding one example *
bonding model element/alloys * physical properties of
There are lots of dot and cross
diagrams i.e. Lewis diagrams of bonding situations
TO CHEMICAL BONDING
Why is knowledge of chemical bonding
Chemists can use the theory of structure
and bonding to explain the physical and chemical properties of materials of
widely varying composition e.g. salt crystals, metals, polymer plastics etc.
etc. Detailed analysis of structures by a variety of techniques shows how
atoms can be arranged in all sorts of ways summarised below with links to
more detailed notes. Chemical bonding theory (covalent, ionic, metallic) explains how atoms are held together in
these different types of structure. This theoretical chemical bonding knowledge,
backed up with experimental evidence, helps
scientists to design and engineer new materials with desirable properties
for specific uses. The properties of these new materials offer new
technological applications and uses in a range of different industrial and
domestic use of technologies from electronic
devices to new structural materials and a lot more besides.
There are three types of strong
chemical bonds: ionic, covalent and metallic. In ionic bonding the particles
(atoms or a group of atoms) form oppositely charged ions. In covalent
bonding the particles are atoms (usually both non-metals) share pairs of
electrons to form the bond. In metallic bonding the metal atoms (actually
positive ions) of the lattice share negative delocalised electrons to bind
themselves together. Ionic bonding occurs in compounds formed from metals
combined with non-metals. Covalent bonding occurs in most non-metallic
elements and in compounds of non-metals. Metallic bonding occurs in metallic
elements and alloys. You should be able to explain chemical bonding in terms
of electrostatic forces and the transfer or sharing of electrons.
When different elements
(different types of atom) react and combine to form a compound (new
substance) chemical bonds must be formed to keep the atoms together. Once these
atoms are joined together its usually difficult to separate them.
The atoms can join together by
sharing electrons in what is known as a covalent bond.
Or, they can transfer or accept
electrons to form positive and negative ions and form an ionic bond.
Metals form another kind of bond
in sharing electrons called a metallic bond.
The types of are briefly explained below with links to even more detailed notes with lots of examples.
Part 1 begins by explaining why
atoms bond together in the first place and then the concepts broadened out
to explain the different types of bonding.
Introduction to some
important definitions in Chemistry
eg atom, molecule, formula, element, compound etc.
are all explained with examples.
Also see notes on How to write
word & symbol equations, work out formula and name compounds formed by
ionic or covalent bonding
Why do atoms bond together?
– 'electron glue'!
Some atoms are very reluctant to combine with other atoms and exist in the air around us as single atoms. These are the Noble Gases and have very stable electron arrangements
e.g. 2, 2,8 and 2,8,8 because their outer shells are full.
The first three are shown in the diagrams
below and explains why Noble Gases are so reluctant to form compounds with other elements.
(atomic number) electron arrangement
All other atoms therefore, bond together to become electronically more
that is to become like Noble Gases in electron arrangement.
Bonding produces new substances and usually involves only the 'outer shell' or 'valency' electrons and atoms can
bond in two ways.
BOND refers to the strong
electrical force of attraction between the atoms or ions in the structure. The
combining power of an atom is sometimes referred to as its valency and
its value is linked to the number of outer electrons of the original uncombined
atom (see examples later).
Each type of chemical bonding is
VERY briefly described below, with links to more detailed notes.
chemical bonding we are dealing with the formation of ions and molecules, so how
big are these particles?
It is difficult to imagine the 'tiny' size of
the particles that we and everything around us is made of.
Atoms and small molecules like water are
around a million times smaller than the width of a human hair!
Most molecules and ions you will come across
in your chemistry studies are 100 000 times smaller than the cells in your body.
comparison data table of particle sizes/dimensions
Examples of dimensions of typical atoms,
molecules and other 'things'!
typical small protein molecule
typical bacteria cell
typical eukaryotic cell
width of a human hair
width in nm
0.3 x 0.6
longest length or diameter in m
1.6 x 10-10
2 x 10-10
2 x 10-10
3.6 x 10-10
6 x 10-10
5 x 10-6
5 x 10-5
~1.0 x 10-4
(a) IONIC BONDING
– an ionic bond is formed by one atom transferring electrons to another atom to form oppositely charged
particles called ions which attract each other – this electrostatic attraction
is called an ionic bond and is most likely formed when a metal combines with a
An ion is an atom
or group of atoms carrying an overall positive or negative
The electric charge is shown as
a superscript +, –, 2+, 2– or 3+ etc.
e.g. Na+, Cl–, [Cu(H2O)6]2+,
If a particle, as in a neutral atom,
has equal numbers of protons (+) and electrons (–) the overall particle charge is zero
i.e. no overall electric charge.
The proton/atomic number in an atom does not change BUT
the number of associated electrons can!
If negative electrons are
excess charge from the protons
produces an overall positive ion.
electrons are gained there is an excess of negative charge, so
a negative ion
The charge on the ion is numerically related to the number of
electrons transferred i.e. electrons lost or gained.
For any atom or group of atoms, for every electron gained
you get a one unit increase in negative charge on the ion, for every electron lost you get
a one unit increase in the positive charge on the ion.
The atom losing electrons forms a
positive ion (cation) and is usually a metallic element.
The atom gaining electrons forms a negative ion
(anion) and is usually a non–metallic element.
The ionic bond then consists of the attractive force between the positive and
negative ions in the structure.
The ionic bonding
forces act in all directions around a particular ion, it is not
directional, as in the case of covalent bonding.
sodium (metal) atom transfers an electron to the chlorine (non–metal) atom
the ionic compound sodium
The bonds between the ions is
very strong and they club together to form a giant ionic lattice with
a very high melting point because it takes a lot of energy to overcome the
attractive forces between the ions - the ionic bonds.
When molten, or dissolved in
water, ionic compounds will conduct electricity because the charged
particles (ions) are free to move and carry the electric current.
For more detailed notes on
this example and lots of other examples ...
(b) COVALENT BONDING
a covalent bond is formed by two atoms sharing electrons so that the atoms
combine to form molecules.
The bond is usually
formed between two non–metallic elements combine to form a molecular compound. The two positive nuclei (due to the positive protons
of both atoms are mutually attracted to the shared negative electrons between
them forming the covalent bond in the molecule. They share the electrons in a way that gives a
stable Noble Gas electron arrangement like helium (2) or neon (2.8) etc..
This kind of bond or
electronic linkage does act in a particular direction i.e. along the
'line' between the two nuclei of the atoms bonded together, this is why
molecules have a particular shape.
and oxygen atoms share electrons to give covalent O–H bonds to form
molecules of the covalent compound water
has a 'bent' shape
METALLIC BONDING isn't
quite like ionic or covalent bonding, although the metal atoms form positive
negative ion is formed from the same metal atoms, but the immobile positive metal
ions/atoms in the lattice are attracted together by the free moving negative electrons between
them. So, like ionic bonding, you do get
attraction between positive and negative particles and this is the metallic
INTERMOLECULAR FORCES – INTERMOLECULAR
Between all particles, but with
particular reference to covalently bonded molecules, there always
exists some very weak electrical attractive forces known as
intermolecular forces or intermolecular bonding.
These constantly acting attractive
forces or intermolecular bonds are very much weaker than full covalent or ionic
chemical bonds (approximately 1/30 to 1/20th
in comparative attractive force).
For example, although the oxygen and
hydrogen atoms are very strongly bonded in water to make a VERY stable
molecule, BUT this does NOT account for
the existence of liquid water and ice!
It is the weak intermolecular forces
that induces condensation below 100oC and freezing–solidification
to form ice crystals below 0oC.
In the reverse process,
when ice is warmed, the
intermolecular forces are weakened and at 0oC the intermolecular
bonds are weakened enough to allow melting to take place.
(evaporation), and particularly at 100oC (boiling), the
intermolecular forces are weak enough for 'intact water molecules'
to escape from the surface of the liquid water.
It is VERY important
to realise that the chemical hydrogen–oxygen covalent bonds
(O–H) in water are NOT broken and the state changes ...
freezing/melting ==> liquid <== condensing/boiling ==>
are due to the
weakening of the intermolecular forces/bonds with increase in
temperature OR the strengthening of the intermolecular
bonds/forces decrease in temperature.
For more details see
Covalent Bonding – small simple molecules and properties
and for Advanced A Level chemistry students:
Introduction to intermolecular forces -
WHY DO SOME ATOM DO NOT READILY FORM
As explained at the start of
Part 1, NOBLE
GASES are very reluctant to
share, gain or lose electrons to form a chemical bond ie they do NOT readily
form a covalent or ionic bond with other atoms.
Noble gases are already
electronically very stable because of their particular electron arrangement.
e.g. 2, 2.8 and
For most other elements
the types of bonding and the resulting properties of the elements or compounds are described in detail
in Parts 2 to 5.
In some of the electronic diagrams ONLY the outer electrons are
Can we deduce the likely chemical
bonding in a material from its physical and chemical properties?
The answer quite simply is YES (in most
as long as you have studied parts 2 to 5 before attempting this question!
The table below describes the
properties of ? compounds. The data is not specific to a substance, just
electrical conduction in solid
electrical conduction when liquid
solubility and electrical conduction in water
Solubility in organic solvents like hexane
decomposes at high temperature
decomposes at high temperature
Can you work out the bonding
of the substance in each case?
Solubility can be a bit
subtle, so take care!
Answers near the bottom
of the page!
PRACTICAL RESEARCH: You can learn how
to classify different types of elements and compounds by investigating their
melting points and boiling points, solubility in water and electrical
conductivity (as solids and in solution) of substances such as sodium chloride,
magnesium sulphate, hexane, liquid paraffin, silicon(IV) oxide, copper sulphate,
and sucrose (sugar). You do simple experiments as well as looking up their
properties in data books.
New bonds formed!
Poetry in motion!
Lots of energy released when metals like
magnesium bond with oxygen!
Ionic Bonding Poem – a snippet of
(anon Y11 student, Whitby
Community College, Oct 31st 2002)
How do I long for a full
being chlorine having
seven, is a horrid hell
but my name is sodium and
I have one spare!
I want to lose it, can we
No? for are we not a
chuck it to me, I promise
then we can live our
and live with full shells
to the end of our days!
and so our tale comes to
as positive and negative
we shall remain friends
Its a good idea to have some idea
of where the elements are in the periodic table, and their electronic structure,
before looking at the theoretical electronic models for ionic, covalent or
The black zig–zag line 'roughly' divides the metals
on the left from the non–metals on the right of the elements of the Periodic
The electronic structures of the first 20
elements of the Periodic Table
You need to know about these to understand the
details of chemical bonding whether it is ionic or covalent etc.
Part of the modern Periodic Table
Pd = period,
Gp = group
metals => non–metals
that H does not readily fit into any group
Chemical Symbol eg 4Be
1 Alkali Metals
Gp 2 Alkaline Earth Metals
Gp 7 Halogens
Gp 0 Noble Gases
Each page has bonding comments about
selected elements highlighted in white
e.g. the type of chemical bond an
element forms with another element (or with itself)
Granddaughter Baby Niamh at nearly 6 months
– first experiment in molecular modelling?
No teething dribbling on the structure please! The greatest chemistry of all
– the chemistry of life!
Answers to the 'type of bonding' question
is an ionic structure and bonding, giant ionic lattice, high
melting/boiling point, only conducts when molten, the solubility and electrical
conduction in water is extra evidence, but isn't definitive for substance A (see
Typical examples would be salts
like sodium chloride, magnesium sulfate
is a giant metallic lattice structure and metal bonds, high
melting/boiling points, high density, conducts in solid, not just liquid.
Typical examples would be iron
simple molecular structure, small molecules with
covalent bonds, low melting/boiling point, no electrical conduction at all
Typical examples would organic
compounds like waxes which dissolve in organic solvents like hexane or
giant covalent lattice, very high melting/boiling, no electrical conduction,
won't dissolve in anything.
Typical examples are carbon
(diamond), silicon, silicon dioxide (silica) and many other minerals found
I've made A to D quite
straightforward (as long as Bonding Parts 2 to 5 have been studied), but I've
simple molecular structure, small molecules with covalent bonds, low
melting/boiling point, no electrical conduction when molten, however it does
conduct when dissolved in water, so ions must be formed to conduct
electricity. The latter is a 'red herring', if it had an ionic structure the
melting/boiling points would be much higher and the liquid would have
Examples are the gases hydrogen
fluoride, hydrogen chloride, hydrogen bromide and hydrogen.
These are all simple molecular
substances, diatomic covalent molecules HF, HCl, HBr, HI), BUT they dissolve
in water to form acids solutions containing the hydrogen ion (H+)
and the corresponding halide ion (F–, Cl–, Br–,
I–), hence the reason why the aqueous solution conducts
electricity. Small covalent molecules often dissolve in organic solvents.
Probably a thermally very stable giant covalent structure, but with weakly
electrical conducting properties (even in the solid) due to delocalised
electrons, completely insoluble. Its unlikely to be an ionic structure
because it conducts in the solid. Metals do not decompose on heating to a
high temperature and all metals will boil.
Examples might be carbon
(graphite) and nanomaterials derived from carbon.
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