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CHEMICAL BONDING Part 1 Introduction to Chemical Bond Formation

Doc 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.

(c) doc b(c) doc bPart 1 Introduction – why do atoms bond together? (this page, read first)

and sub–index for Parts 2–5 (this page)

Part 2 Ionic Bonding – compounds and properties

Part 3 Covalent Bonding – small simple molecules and their properties

Part 4 Covalent Bonding – macromolecules and giant covalent structures

Part 5 Metallic Bonding – structure and properties of metals

Extra notes on chemical bonding for ADVANCED A Level Students ONLY (IB, US grade 11-12)

6.1 Electronegativity, bond polarity, type of chemical bonding

6.2 More on ionic structures and ionic bonding

(Working out 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

Introduction - VSEPR theory - lots of inorganic molecule/ion examples

Some other molecules/ions of carbon, nitrogen, sulphur and chlorine

shapes and bond angles of organic molecules

Intermolecular forces - intermolecular bonding

Introduction to intermolecular forces - intermolecular bonding

Detailed comparative discussion of boiling points of 8 organic molecules

Boiling point plots for six organic 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 molecules

NANOSCIENCE – NANOTECHNOLOGY – NANOCHEMISTRY (index of pages and keyword index)

SMART MATERIALS 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 combination), magnesium/calcium 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 * polymers/plastics * properties of polymers * inter/intra (internal)–molecular forces * hydrogen H2, chlorine Cl2, hydrogen chloride HCl, water H2O, ammonia NH3, methane CH4, oxygen O2, carbon dioxide CO2, ethene C2H4, nitrogen N2, ethane C2H6, chloromethane CH3Cl, methanol CH3OH, carbon (diamond), carbon (graphite), carbon (buckminsterfullerene/fullerenes), silica/silicon dioxide SiO2 * examples of ionic compounds * physical 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 * metal bonding model element/alloys * physical properties of metals *

There are lots of dot and cross diagrams i.e. Lewis diagrams of bonding situations


Why is knowledge of chemical bonding important?

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

Part 1. 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.

(c) doc b (c) doc b (c) doc b (atomic number) electron arrangement

All other atoms therefore, bond together to become electronically more stable, 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.

The phrase CHEMICAL 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.


In 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.

A comparison data table of particle sizes/dimensions

Examples of dimensions of typical atoms, molecules and other 'things'!

material carbon atom water molecule sodium ion chloride ion glucose molecule typical small protein molecule typical bacteria cell typical eukaryotic cell width of a  human hair
symbol-formula C H2O Na+ Cl C6H12O6 - - - -
diameter or width in nm 0.16 0.20 0.20 0.36 0.3 x 0.6 ~5 5000 50000 ~100000
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 10-9 5 x 10-6 5 x 10-5 ~1.0 x 10-4
********** ************** *************** ************* *************** ************** ************** *********** ************ ***************


(c) doc b(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 non-metal.

  • An ion is an atom or group of atoms carrying an overall positive or negative electric charge

    • The electric charge is shown as a superscript +, –, 2+, 2– or 3+ etc.

    • e.g. Na+, Cl, [Cu(H2O)6]2+, SO42– etc.

  • 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 lost the excess charge from the protons produces an overall positive ion.

  • If negative electrons are gained there is an excess of negative charge, so a negative ion is formed.

  • 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.

  •  (c) doc b(c) doc bThe sodium (metal) atom transfers an electron to the chlorine (non–metal) atom in forming the ionic compound sodium chloride

  • 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 in them) 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..

(c) METALLIC BONDING isn't quite like ionic or covalent bonding, although the metal atoms form positive ions, no 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 bond.


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.

Above 0oC (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 ...

solid <== freezing/melting ==> liquid <== condensing/boiling ==> gas ...

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 - intermolecular bonding


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 2.8.8 etc.

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 shown.

(f) Can we deduce the likely chemical bonding in a material from its physical and chemical properties?

The answer quite simply is YES (in most cases!), 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 'typical properties'

Substance Melting Point Boiling Point density (g/cm3) electrical conduction in solid electrical conduction when liquid solubility and electrical conduction in water Solubility in organic solvents like hexane
A 900oC 1720oC 1.5 none good soluble, good insoluble
B 1200oC 2000oC 7.5 good good insoluble insoluble
C 65oC 260oC 1.2 none none insoluble soluble
D 1100oC 2500oC 3.7 none none insoluble insoluble
E –120oC –65oC very small none none soluble, good soluble
F decomposes at high temperature decomposes at high temperature 1.2 moderate decomposes insoluble insoluble

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 chemical poetry

(anon Y11 student, Whitby Community College, Oct 31st 2002)

How do I long for a full outer shell!

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 not share?

No? for are we not a perfect match

chuck it to me, I promise to catch

then we can live our separate ways

and live with full shells to the end of our days!

and so our tale comes to an end

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 metallic bonding

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 Table.

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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.

Pd metals Part of the modern Periodic Table

Pd = period, Gp = group

metals => non–metals
Gp1 Gp2 Gp3 Gp4 Gp5 Gp6 Gp7 Gp0

1H  Note that H does not readily fit into any group

2 3Li 4Be atomic number Chemical Symbol eg 4Be 5B 6C 7N 8O 9F 10Ne
3 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar
4 19K 20Ca 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr
5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe
6 55Cs 56Ba Transition Metals 81Tl 82Pb 83Bi 84Po 85At 86Rn
Gp 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

A 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 E).

Typical examples would be salts like sodium chloride, magnesium sulfate

B 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 and copper.

C simple molecular structure, small molecules with covalent bonds, low melting/boiling point, no electrical conduction at all (no ions).

Typical examples would organic compounds like waxes which dissolve in organic solvents like hexane or propanone ('acetone').

D 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 in rocks.

I've made A to D quite straightforward (as long as Bonding Parts 2 to 5 have been studied), but I've

E 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 conducted.

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.

F 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.

WHAT NEXT? and other associated Pages

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