Why do atoms bond ?

As you have learned in the 'Structure of the atom' topic, electrons are placed into energy levels, and the electron arrangement can be given for each atom.
The most stable compounds in the periodic table, are the Noble gases (you learned this in the 'substances' topic). The electron arrangement of these gases all have the maximum number of electrons in their outer shell.

	Helium	Neon	Argon
Electron arrangement :	2	2,8	2,8,8

Scientists believe that when the outer electron shell is full, this is the most stable arrangement of electrons and is called a stable electron arrangement .
When atoms bond to each other, they attempt to achieve the same electron arrangement of the nearest noble gas. They do this in several different ways and only the outer shell electrons are used in forming bonds.

 
The covalent bond

As with all bond formations, the atoms in question must first collide with one another (as mentioned earlier in the 'Reaction Rates' topic).

When some atoms collide with each other, the electrons in the outer shell can be shared between the atoms. Although the electrons of the two atoms are both negatively charged and repel each other, when a collision takes place with sufficient energy to form a compound, the outer energy levels overlap and the atoms share the electrons. The overlap area has an increase in negative charge, which is strongly attracted by the positive nuclei of both atoms. This effectively draws the two atoms closely together and a strong force of attraction between the nuclei and the shared electrons forms a strong covalent bond.

Covalent bonds only occur between non-metal atoms and the bond is a strong force of attraction.

When covalent bonds are formed between a small number of non-metal atoms, the resulting compound is called a molecule.

Some examples showing the sharing of outer shell electrons to obtain a stable electron arrangement are :

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Electron clouds

In the 'Structure of the atom' topic, you learned that electrons are found in the space around the nucleus of the atom. You also learned in the 'Structure of the atom' topic, that an atom's electrons are ordered into energy levels or shells.

Within these shells, the electrons are found in electron clouds or orbitals.

• Electron clouds have different shapes depending on the energy level the electrons are in.
• The shape of the electron cloud defines the area in which its electrons can be found
• Each electron cloud can hold a maximum of 2 electrons
• Electrons fill these clouds / orbitals singly if possible, but they double up after each empty cloud on the energy level contains one electron.
• Electrons normally only move onto the next energy level once the current energy level is full.

Some electron clouds are shown below :

1st energy level	2nd energy level (with 4 or more electrons)	3rd energy level (with 4 or more electrons)
1 electron cloud	4 electron clouds	4 electron clouds
(1x2)=2 electrons	(4x2)= 8 electrons	(4x2)= 8 electrons

For example, a hydrogen atom, which has 1 electron, has the electron arrangement 1.

This electron is in the first energy level which is in the shape of a sphere.

(In the diagrams that follow, electrons are indicated by the small dots inside the orbitals)

For example, an oxygen atom, which has 8 electrons, will have the electron arrangement, 2,6

This means that two electrons are in the inner shell (sphere shaped) and the other six electrons are in the outer shell (in the shape of a tetrahedron). This is shown below.

2D representation of 1st and 2nd energy levels	3D representation of 2nd energy level only

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Polar covalent bonding

When covalent bonds are formed, the electrons that are shared between the two atoms are not always equally shared. This is because some atoms attract electrons more strongly than others. This attraction is called electronegativity.

The only type of covalent bond that has an equal share of electrons between the two atoms is when the two atoms are of the same element. e.g. H₂, O₂, Cl₂. These bonds are called pure covalent bonds.

All other covalent bonds are classed as polar covalent bonds.

For example, the molecule water has two elements : oxygen and hydrogen.
The oxygen atom attracts the electrons more strongly than the hydrogen atoms and as a result, pulls the shared electrons closer to itself. This makes the oxygen atom slightly more negatively charged (δ-) than the hydrogen atoms (δ+). This creates a permanent dipole (meaning two poles or charges on the molecule).

The interactions of permanent dipoles between water molecules

When molecules contain polar covalent bonds, the molecules have a slightly positive end which can attract the negatively charged ends of other molecules and vice-versa.

Molecules that have polar covalent bonding are classed as polar molecules.

This type of attraction is weak compared to the covalent bond.

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Discrete covalent substances

Discrete covalent substances are made up of molecules.

When a molecule contains only two atoms (e.g. H₂, Cl₂, HCl), it is called a diatomic molecule.

All of the Halogens (group 7) exist as diatomic molecules, as do Hydrogen, Nitrogen and Oxygen.

When a molecule contains only three atoms (e.g. CO₂, H₂O, O₃), it is called a triatomic molecule.

All discrete covalent substances (with polar and non-polar bonding) exhibit a type of attraction between the molecules. This attraction is called van der Waals forces.

The interactions of temporary dipoles between molecules

This attraction arises when random electron movement within the orbitals causes one side of a molecule to have a slightly negatively charged end and a slightly positively charged end (a temporary dipole). These charges cause molecules to attract one another in the same way that a polar molecule does. This force of attraction is much smaller than that in a covalent bond.

Because these charges (or dipoles) are randomly distributed and are not permanent, this type of van der Waals force is usually smaller the attraction between polar molecules.

 
Multiple covalent bonds

Sometimes an atoms can form more than one bond with each other in an attempt to achieve a stable electron arrangement.

e.g. Oxygen has an electron arrangement of 2,6 and requires a share of a further 2 electrons to reach the electron arrangement of Neon (2,8).

Oxygen exists as a diatomic element, i.e. it forms covalent bonds with itself.
Since each oxygen atom is short of two electrons, it can form two covalent bonds with the other atom. This is called a double bond and involves the sharing of 4 electrons between two atoms.

A double bond between oxygen atoms forming a molecule of oxygen

e.g. Nitrogen also exists as a diatomic element. This involves two nitrogen atoms sharing 6 electrons with each other (three from each atom) forming a triple bond.

A triple bond between nitrogen atoms forming a molecule of nitrogen

Elements that require four electrons to achieve a stable electron arrangement do not form 4 bonds with each other as there is too much strain involved in the bond. Instead they can form a combination of single, double and triple bonds to gain a stable arrangement.

 
The shapes of molecules

The shapes of molecules are based on the shapes of the orbitals that are used in bonding.

When a carbon atom (electron arrangement - 2,4) bonds to four hydrogen atoms to form a compound called methane, or carbon hydride, the carbon shares an electron with each of the four hydrogen atoms to achieve the same electron arrangement as Neon (2,8), and the four hydrogen atoms also achieve a stable arrangement of Helium (2) as shown on the left.

Since carbon's bonding electrons are in the second energy level (2,4), they are arranged in the shape of a tetrahedron as shown on the right.

When this atom of carbon bonds to four hydrogen atoms the molecule has the shape of a tetrahedron and looks like this:

 
When a nitrogen atom bonds (covalent bonds) with three hydrogen atoms (to form a molecule of ammonia), the resulting molecule shape is called a pyramidal structure and is shown below :

Notice that there is a pairs of electrons that are not involved in the formation of the three single covalent bonds. These electrons are called lone pairs, and they take up one of the tetrahedral positions and affect the shape of the molecule.

 
When an oxygen atom bonds (covalent bonds) with two hydrogen atoms (to form a molecule of water), the resulting molecule shape is called a bent structure and is shown below (the oxygen atom has two lone pairs of electrons that are not involved in bonding) :

 
Chemical formulae and structural formulae

The chemical formula tells us the elements that are in a compound and it also tells us the number of each type of atom that is present.

For example, the compound hydrogen oxide (or water) contains hydrogen and oxygen only (you learned this in the 'Substances' topic).

You also know from the previous examples that oxygen needs to share an electron with two hydrogen atoms to achieve a stable electron arrangement.
This means that there are two hydrogen atoms for every one oxygen atom.

This gives the chemical formula of hydrogen oxide / water as H₂O.

Notice that the number '1' (showing one atom of oxygen) is not written in the chemical formula.

The full structural formula of a compound also tells us the elements present and the number of atoms of each, but in addition tells us the way that these atoms are arranged in the compound.

The chemical formulae and full structural formulae of some other compounds are shown below :

Chemical name	Atoms present	Chemical Formula	Full Structural Formula
Hydrogen oxide (water)	H & O	H2O
Nitrogen hydride (ammonia)	N & H	NH3
Carbon hydride (methane)	C & H	CH4
Ethanol (alcohol)	C, H & O	C2H5OH

 
Covalent networks

As well as discrete covalent compounds and small molecules, bonds can be formed between many atoms forming a giant compound containing many hundreds, thousands, or millions of atoms.

These compounds are called covalent networks and consist of a giant lattice of covalently bonded atoms.

For example, the element carbon can exist in the form of diamond (as you learned in the 'Substances' topic), which is a covalent network structure.

The full structural formula of diamond is nearly impossible to write as it contains a huge number of carbon atoms.

In the same way, the compounds silicon carbide and silicon oxide form giant lattices / covalent networks. Part of the structures of these are shown below :

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Since these substances contain a huge number of atoms, it is not practical to write the chemical formula for them. Instead we write the empirical formula .

The empirical formula shows us the simplest ratio of elements in the substance.

For example :
The empirical formula of silicon carbide (1 silicon atom for every 1 carbon atom) is SiC
The empirical formula of silicon oxide (1 silicon atom for every 2 oxygen atoms) is SiO₂

For a covalent network, the chemical formula is written as the empirical formula.

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Ionic bonding

An ionic bond (sometimes called an electrovalent bond) usually occurs between a metal and a non-metal and involves ions, which are charged atoms (or groups of atoms).

In ionic bonding, electrons are donated from one atom to another allowing both atoms to achieve a stable electron arrangement.

If we consider the group 1 metals, they have only have one electron in their outer shell.
In order to achieve a stable arrangement, these metals 'lose' their single outer shell electron, leaving the next lower energy level full.

For example, the element sodium (electron arrangement 2,8,1) can lose its outer shell electron to achieve the electron arrangement 2,8. This leaves a 1+ (one positive) charge on the atom and this is called a positive ion.

Some metals (group 2) have 2 electrons in their outer shell and lose these to gain a stable electron arrangement. When this occurs they have a charge of 2+.

Similary aluminium can lose 3 outer shell electrons to achieve a stable electron arrangement leaving a 3+ charge on the atom.

Some examples of positive ions are shown in the table below.

In the same way, non-metal atoms can also gain electrons from the metal in order to reach a stable electron arrangement. For example, chlorine which has an electron arrangement of 2,8,7 requires one electron to fill its outer shell, making it 2,8,8. This in turn gives the chlorine a negative charge as it has one more electron than protons in the nucleus.
This is called a negative ion.

Some examples of negative ions are shown in the table below.

When a metal gives its outer shell electron(s) to a non-metal, the positive ion and negative ions that are formed attract one another and form an ionic bond.

For example, sodium and chlorine atoms would form an ionic bond making the compound sodium chloride as shown below :

There are also ions that exist that contain more than one element. These are called group ions and some examples are the sulphate ion (SO₄^2-) and the ammonium ion (NH₄⁺). The formulae for these group ions are found in the data booklet.

A small portion of the Sodium chloride lattice The chemical formula is NaCl

Ionic compounds often form lattices which are called ionic structures.
These structures are held together due to attraction between neighbouring, oppositely-charged ions.

As with covalent network structures, the empirical formula is given rather than the chemical formula.

Ionic bonds are extremely strong.

 
Metallic bonding

Metals, as mentioned in ionic bonding, can lose their outer shell electrons to gain a stable electron arrangement. Metal atoms are arranged in a lattice and can also delocalise their outer shell electrons, allowing them to move freely between the atoms in the lattice.

This effectively brings the metals closer to obtaining a stable electron arrangement and since the metal atoms become positively charged (ions) they attract the free moving electrons in the lattice. This attraction forms a metallic bond which is very strong.

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Properties of substances due to type of bonding

Electrical conductivity

A substance will conduct electricity when charged particles can move through the material. When this happens, the material is said to be an electrical conductor, and if the material does not allow the movement of charged particles through it, it is called and insulator.

Charged particles can be electrons or ions. When a material conducts electricity by allowing electrons to move through it, the material is not altered chemically. However, when a substance conducts electricity by allowing ions to move through it, these substances are usually broken up in the process (this is called electrolysis and is mentioned further down the page).

Covalent compounds do not conduct electricity as they do not contain any charged particles.

Ionic compounds contain ions and will conduct electricity when these particles can move freely through the material. This only occurs when the ionic compound is in solution or when it is molten (has been melted into a liquid). An ionic compound in solution or as a melt is called and electrolye. Solid ionic compounds do not conduct electricity.

All metallic elements conduct electricity as they have electrons moving freely through their structures.

Carbon in the form of graphite also has delocalised electrons in its structure and as a result will conduct electricity.

You can test the electrical conductivity of some elements in the Virtual Lab.

 
Melting points and boiling points

To melt a substance, the bonding between units (molecules and ionic structures) has to be substantially weakened (or partly broken). To boil a substance, these bonds have to be broken.

Note that the covalent bonds between atoms are not broken during melting or boiling of discrete molecular substances such as H₂, only the attraction between molecules and units - these forces of attraction are called van der Waals forces (present in all substances), those between polar molecules (covalent substances) and those between neighbouring ions (ionic compounds). However when a covalent network solid such as graphite or silicon carbide is melted, covalent bonds have to be broken.

• Discrete covalent substances have low melting and boiling points eg N₂, O₂.

• Almost all substances (except mercury) that exists as gases or liquids at room temperature, are discrete covalent substances.

• Polar covalent substances have higher melting and boiling points than pure covalent substances due to the increased attraction between molecules.

• Covalent networks have high melting and boiling points and are all solids at room temperature. This is because the structure is held together by strong covalent bonds which have to be broken to allow melting.

• Ionic compounds have very high melting and boiling points and are all solids at room temperature. This is because ionic bonds need to be weakened/broken to melt or boil them.

• Metallic elements have high melting and boiling points. This is because metallic bonds need to be weakened/broken to melt or boil them.

 
Colours of ionic compounds

The colours of ionic substances depends on the colour of the ions present.

The colours of some ions are listed below :

For example, sodium chloride is likely to be colourless, copper sulphate is likely to be blue, potassium chromate is likely to be yellow and copper chromate is likely to be green.

 
Solubility

Most ionic substance dissolve in water and this involves the lattice being broken up completely.

Some covalent substances dissolve in water, but most dissolve in other solvents.

Usually the rule 'like dissolves like' applies - This means that :

• non-polar solutes will dissolve in non-polar solvents
• polar solutes will dissolve in polar solvents
• ionic solutes will dissolve in polar solvents

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Electrolysis

Electrolysis is the use of electricity to break up an ionic compound either in solution or as a melt (an electrolyte).

Covalent compounds or metallic elements cannot undergo electrolysis.

Normally, a d.c. supply is used during electrolysis. One electrode is connected to the positive end of the supply and is called the anode. The other electrode is connected to the negative end of the supply and is called the cathode.

During electrolysis, the positive ions are attracted to the negative electrode and the negative ions are attracted to the positive electrode.

At the positive electrode, negatively charged non-metal ions lose electrons.

At the negative electrode, positively charged ions gain electrons.

For example, when a solution of copper chloride undergoes electrolysis, the copper ions are attracted to the negative electrode and the chloride ions are attracted to the positive electrode.

The copper ions gain two electrons and form copper metal, and the chloride ions lose two electrons to form chlorine gas.

Cu²⁺(aq) + 2e^- Cu(s)
2Cl^-(aq) Cl₂(g) + 2e^-

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New words and their meanings

Stable electron arrangement - An atom is most stable when it has the electron arrangement of the nearest noble gas in the periodic table.

Covalent bond - Formed when atoms share electrons to achieve a stable electron arrangement.

Electron cloud / Orbital - The space that an electron occupies in an atom. Each energy level has different electron clouds associated with it.

Pure covalent bond - A covalent bond formed between two atoms of the same element. This means the atoms have an equal share of the electrons in the bond.

Polar covalent bond - A covalent bond formed between two atoms in which the atoms have an unequal share of the bonding electrons. The atom which has the greatest share has as δ- charge and the atom with the least share has a δ+ charge associated with it.

Electronegativity - The strength with which an atom attracts electrons.

Discrete covalent substances - These are made up of small molecules.

Van der Waals forces - The forces of attraction between molecules (and atoms) caused by temporary positive and negative charges being formed due to random electron movement.

Dipole - Formed when a molecule (or atom) has a positively charged end and a negatively charged end. These can be permanent (as in polar molecules) or temporary (as in van der Waals forces caused by random electron movement).

Lone pair - Two electrons in an orbital that are not used in bonding.

Chemical formula - Tells us the atoms present and the quantity of each in a compound.

Full structural formula - Tells us the atoms present, the quantity of each and their arrangement in a compound.

Covalent network - A giant lattice of covalently bonded atoms.

Empirical formula - A type of chemical formula that gives the simplest ratio (not the actual numbers) of atoms in a compound.

Ionic bond - A bond that involves the transfer of electrons from one particle to another to form positively and negatively charged particles. These particles attract one another and form an ionic bond.

Ions - A charged atom (or group of atoms) : A positive ion is an atom (or group of atoms) that has lost electrons and a negative ion is an atom (or group of atoms) that has gained electrons.

Metallic bond - A bond formed between metal atoms, where the outer shell electrons are delocalised and allowed to move freely throughout the lattice of metal atoms.

Electrolysis - The breaking down of a compound using electricity.