Chemistry AS - Chapter 03 - Bonding

Revision cards for Chapter 3- Bonding

The Nature of Ionic Bonding

Bonds between atoms always involve their outer electrons.

Noble gases have an outer shell full of electrons and are unreactive.

When atoms bond, they share or transfer electrons to achieve a more stable electron arrangement (a full outer shell).

Three types of strong chemical bonds:

  • Ionic
  • Covalent
  • Metallic

Metals have one, two, or three electrons in their outer main levels.

The easiest way for them to attain a full outer shell is to lose their outer electrons.

Non metals have space in their outer shell. The easiest way for them to attain a full outer shell is to gain electrons.

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The Nature of Ionic Bonding

Ionic bonding occurs between metals and non metals.

Electrons are transferred from metal atoms to non metal atoms.

Positive and negative ions are formed.

Sodium (Na) has 11 electrons. Configuration: 2,8,1

Chlorine (Cl) has 17 electrons. Configuration: 2,8,7

An electron is transferred from Na to Cl.

Each outer shell is full.

Na becomes positively charged as it has lost an electron.

Cl becomes negatively charged as it has gained an electron.

The two ions are attracted to each other by electrostatic forces.

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The Nature of Ionic Bonding

This electrostatic attraction extends throughout the compound.

Every positive ion attracts to a negative ion.

The structure is a lattice. This is because there is also repulsion occuring between similarly charged ions.

Ionic compounds are always solid at room temperature.

They have giant structures and therefore high melting points.

This is due to the amount of energy needed to break up the lattice of ions and the electrostatic attraction between ions.

Ionic compounds conduct electricity when molten or dissolved. NOT when they are solid. Ions that carry a current are free to move in a liquid state.

Ionic compounds are brittle and shatter easily.

A blow could produce contact between like charges and break the structure.

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Covalent Bonding

A covalent bond forms between a pair of non-metal atoms.

Atoms share some of their outer electrons so that each atom has a full outer shell.

A covalent bond is a shared pair of electrons.

Atoms with covalent bonds are held together by the electrostatic attraction between the nuclei and the shared electrons.

In a double covalent bond, the four electrons are shared .Example: Oxygen.

The two atoms in the oxygen molecule share two pairs of electrons.

Substances composed of molecules are gases, liquids or solids with low melting temperatures.

This is because the covalent bonds work only between the atoms within the molecule. There is a weak attraction between the molecules. Therefore less energy is needed to move them.

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Covalent Bonding

They are poor conductors of electricity because the molecules are overall neutral.

If they dissolve in water, they do not conduct electricity.

A single covalent bond consists of a pair of electrons shared between two atoms.

In most covalent bonds, each atom provides one of the electrons. However, in some bonds, one atom provides both electrons.

This is called co-ordinate bonding (dative covalent bonding).

In a co-ordinate/dative covalent bond:

  • The atom which is accepting the electron pair is an atom that does not have a filled outer main level of electrons. It is electron deficient.
  • Atom that is "donating" has a pair of electrons not being used for a bond. This is called a "lone pair". (Ammonium ion is an example of this)
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Covalent Bonding


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Covalent Bonding

Co-ordinate bonds are represented by an arrow -->

The arrow would point towards the atom that is accepting the electron pair. In the Ammonium Ion, the arrow would point from N to the H on the right of it.

Ammonium ion is completely symmetrical and all the bonds have exactly the same strength and length.

Ammonium ion has covalently bonded atoms but it is a charged particle.

To spot whether it is an ionic bond or a co-ordinate bond in a charged molecule, just check what type of atoms they are.

If it is metal and non-metal, it is ionic. If it is non-metal and non-metal, it is co-ordinate.

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Electronegativity - Bond Polarity in Covalent Bond

Electronegativity is the power of an atom to attract the electron density in a covalent bond towards itself.

Fluorine is better than hydrogen at attracting electrons. Therefore we say Fluorine is more electronegative than hydrogen.

The Pauling Scale is used to measure electronegativity.

Electronegativity depends on:

  • Nuclear Charge
  • Distance between the nucleus and the outer shell electrons.
  • The shielding of nuclear charge by electrons in inner shells.

The smaller the atom, the closer the nucleus is to the shared outer main level electron. It's electronegativity would be greater.

The larger the nuclear charge, the greater the electronegativity.

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Electronegativity - Bond Polarity in Covalent Bond

As you go up a group, electronegativity increases as shielding decreases and atoms get smaller, so the nuclear charge is closer to the outer shell.

Across a period, electronegativity increases. As nuclear charge increases, the number of main levels remain the same and atoms get smaller.

Most electronegative atoms are found on the top right hand corner of the periodic table.

Polarity is about the unequal sharing of the electrons between atoms that are bonded covalently. It is a property of the covalent bond.

When both atoms are the same, the electrons must be shared equally between the atoms.

Between 2 atoms with different electronegativity, the electrons in the bond will not be shared equally between the atoms.

Eg. HF. The electrons are more attracted towards F than H.

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Electronegativity - Bond Polarity in Covalent Bond

Therefore, we add partial charges. We add a delta on top of each atom and put a + or - accordingly.

A - would be put above the more electronegative atom. A + would be put on the less electronegative atom.

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Metallic Bonding

Metals are shiny elements made up of atoms which can easily lose up to three electrons, forming positive metal ions.

Atoms in a metal element cannot transfer electrons (as in ionic bonding) unless there is a non-metal atom present to recieve them. In a metal element, the outer main levels of the atom merge.

The outer electrons are no longer associated with any one particular atom.

Metallic bonding is where a lattice of positive ions exist in a "sea" of outer electrons which are delocalised. Delocalised means they are not tied down to any particular atom.

Positive ions tend to repel one another but this is balanced by the electrostatic attraction of these positive ions for the negatively charged "sea" of delocalised electrons.

The number of delocalised electrons depends on how many electrons lost by each metal atom. Metallic bonding spreads across the metal (giant structure).

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Metallic Bonding

Metals are good conductors of electricity and heat.

Electricity - an electron from the negative terminal of the supply joins the sea of delocalised electrons at the same time when a different electron leaves from the positive terminal.

Heat - metals have high thermal conductivities. The sea of electrons is partly responsible for this property, with energy also spread by the increasingly vigorous vibrations of the closely packed ions.

The strength of the metallic bonding depends on:

  • The charge of the ion (the higher the charge, the greater number of delocalised electrons and therefore the stronger the electrostatic attraction between the positive ions and the negative sea of electrons)
  • The size of the ion (the smaller the ion, the closer the electrons are to the positive nucleus and the stronger the bond)

Metals tend to be strong because of the delocalised electrons extending through the structure. No individual bonds to break.

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Metallic Bonding

Metals are malleable and ductile.

Malleable means they can be beaten into shape.

Ductile means they can be stretched and pulled into shape.

After a small distortion, each metal ion is in exactly the same environment as before. Therefore the new shape is retained. This contrasts with the brittleness of ionic compounds.

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Metallic Bonding

Metals generally have high melting and boiling points.

This is because they have giant structures.

There is strong attraction between the metal ions and the delocalised sea of electrons. 

This makes the atoms difficult to seperate.

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Forces Acting Between Molecules

Atoms in molecules are held together by strong covalent, ionic or metallic bonds. Molecules and seperate atoms are attracted to one another through weaker forces called intermolecular forces. If the intermolecular forces are strong enough, the molecules could form a liquid or solid.

Three types:

  • van der Waals forces - acts between all atoms and molecules.
  • Dipole-Dipole forces - acts only between certain types of molecules.
  • Hydrogen bonding - acts only between certain types of molecules.

van der Waals forces are the weakest, Dipole-Dipole are stronger, Hydrogen bonding are the strongest.

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Forces Acting Between Molecules

Molecules with polar bonds may have a dipole moment. This sums up the effect of the polarity of all the bonds in the molecule.

In molecules with more than one polar bond, the effects may cancel, to leave a molecule with no dipole moment. The effects may also add up and reinforce each other. It depends on the shape of the molecule.

Carbon Dioxide is a linear molecule and the dipoles cancel.

Tetrachloromethane is tetrahedral and the dipoles also cancel out.

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Forces Acting Between Molecules


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Forces Acting Between Molecules

In dicholoromethane, the dipoles do not cancel out due to the shape of the molecule. 

Dipole-Dipole forces act between molecules that have permanent dipoles. Example: hydrogen chloride. Chloride is more electronegative than hydrogen, so the electrons are pulled towards the chlorine atom rather than the hydrogen atom.

The molecule has a dipole and is written H(delta+) - Cl(delta -)

Two molecules which have dipoles will attract one another.

 Whatever their starting positions, the molecules with dipoles will "flip" to give an arrangement where the two molecules attract.

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Forces Acting Between Molecules

All atoms and molecules are made up of positive and negative charges, even though they are neutral overall.

These charges produce very weak electrostatic attractions between all atoms and molecules. We call these van der Waals forces.

Take the helium atom. It has 2 positive charges in its nucleus (2 protons) and 2 negative charges surrounding the nucleus (2 electrons).

The atom as a whole is neutral but at any moment in time the electrons could be anywhere. This means the distribution of charge is changing at every instant.

This means that the atom has a dipole at that moment. An instant later, the dipole may be in a different direction. But the atom will have a dipole at any point in time even if it is a temporary dipole.

This dipole affects the electron distribution in nearby atoms so that they are attracted to the original helium atom for that instant.

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Forces Acting Between Molecules

As the electron distribution of the original atom changes, it will induce new dipoles in the atoms around it which will be attracted to the original one. These are van der Waal forces.

van der Waal forces act between all atoms or molecules at all times.

They are in addition to any other intermolecular forces.

The dipole is caused by the changing position of the electron cloud, so the more electrons there are, the larger the instantaenous dipole will be.

The size of van der Waals forces increases with the number of electrons present. This means that atoms or molecules with large atomic/molecular masses produce stronger van der Waals forces.

This explains why the boiling points of noble gases increase as the atomic number increases and why the boiling points of hydrocarbons increase with increased chain length.

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Hydrogen Bonding

Hydrogen bonding has charcteristics of dipole-dipole attraction and some of a covalent bond. It consists of a hydrogen atom "sandwiched" between two very electronegative atoms.

There are conditions that have to be present in order for hydrogen bonding to occur.

We need a very electronegative atom with a lone pair of electrons covalently bonded to a hydrogen atom.

Water molecules fulfil these condiitons. Oxygen is much more electronegative than hydrogen so water is polar.

We would expect to find weak dipole-dipole attractions between the molecules but in this case, the intermolecular bonding is much stronger for two reasons:

1. Oxygen atoms in water have lone pairs of electrons.

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Hydrogen Bonding

2. Hydrogen atoms are highly electron deficient in water. This is because oxygen is very electronegative and attracts the shared electrons in the bond towards it. The hydrogen atoms in water are positively charged and very small.

The lone pair of electrons on the oxygen atom of another water molecule is strongly attracted to the electron deficient hydrogen atom.

This strong intermolecular force is called a hydrogen bond. Hydrogen bonds are considerably stronger than dipole-dipole attractions, though much weaker than a covalent bond.

Represented by dashes "---"

The only atoms that are electronegative enough to form hydrogen bonds are oxygen, nitrogen and fluorine.

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Hydrogen Bonding

The effect of hydrogen bonding between molecules can be seen if we look at the boiling points of hydrides of elements of Group 4,5,6 and 7 plotted against period number.

The noble gases show a gradual increase in boiling point because the only forces acting between the atoms are van der Waals forces and these increase with the number of electrons present.

The boiling points of H2O, HF and NH3 are higher than those of the hydrides of other elements in their group. This is because hydrogen bonding is present in these compounds and these are stronger intermolecular forces.

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States of Matter

When we heat a solid and supply energy to the particles, it makes them vibrate more about a fixed position. This slightly increases the average distance between particles so the solid expands.

In order to turn a solid into a liquid, we need to supply more energy. This energy is needed to weaken the forces that act between the particles which hold them in the solid state. The energy needed is called the enthalpy change of fusion.

When we heat a liquid, we supply energy to the particles which make them move more quickly. We say they have more kinetic energy. On average, the particles move a little further apart so liquids also expand on heating.

In order to turn a liquid into a gas, we need to supply enough energy to break all the intermolecular forces between the particles. A gas consists of particles that are far apart and moving independently. The energy needed is called the enthalpy change of vaporisation.

When we heat a gas, the particles gain kinetic energy and move faster. They get much further apart so gases expand when heated.

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States of Matter

Crystals are solids where the particles have a regular arrangement and are held together by forces of attraction. These could be strong bonds (covalent, ionic, metallic) or weaker intermolecular forces (van der Waals, dipole-dipole or hydrogen).

The strength of the forces of attraction between particles in the crystal affects the physical properties of the crystals.

Stronger the force, the higher the melting temperature and greater the enthalpy of fusion.

There are four basic crystal types: ionic, metallic, molecular and macromolecular.

Ionic crystals - strong electrostatic attractions between oppositely charged ions. Example: sodium chloride. High melting points due to strong electrostatic attractions throughout the structure. Requires a large amount of energy to break.

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States of Matter

Metallic Crystals - Exist as a lattice of positive ions embedded in a sea of delocalised electrons. Attraction of positive to negative extends throughout the structure. High melting temperature due to strong metallic bonds.

Molecular Crystals - Molecules are held in a regular array by one or more of the three intermolecular forces (van der Waals, dipole-dipole and hydrogen bonding).

Covalent bonds within the molecules hold the atoms together but they do not act between the molecules. Intermolecular forces are much weaker than the covalent, ionic or metallic bonds.

Molecular crystals have a low melting point and low enthalpies of fusion.

Iodine is a molecular crystal. Strong covalent bonds hold pairs of iodine atoms together to form I2 molecules. Since they have a large number of electrons, the van der Waals forces are strong enough to hold the molecules together as a solid.

However, van der Waals forces are much weaker than covalent bonds.

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States of Matter

Iodine crystals are:

  • soft and break easily
  • Low melting temperature and sublimes readily to form gaseous iodine molecules.
  • Cannot conduct electricity because there are no charged particles to carry charge.

Macromolecular crystals - covalent compounds are not always made up of small molecules. In some substances, covalent bonds extend throughout the compound and we have the typical property of a giant structure held together with strong bonds: high melting temperature.

Diamond and Graphite are made up of carbon only. They are polymorphs or allotropes of carbon as they are very different materials because they are arranged/bonded differently. They are macromolecular structures.

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States of Matter

Diamond - pure carbon with covalent bonding between every carbon atom. The bonds spread throughout the structure which is why it is a giant structure.

Carbon atom has 4 electrons in its outer shell. In diamond, each atom forms 4 single covalent bonds with other carbon atoms. The four electron pairs repel each other.

The atoms form a giant three-dimensional lattice of strong covalent bonds. Diamond has the following properties:

  • Very hard material
  • Very high melting point (3700K)
  • Does not conduct electricity as there are no free charged particles to carry charge.
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States of Matter

Graphite - Consists of pure carbon but atoms are bonded and arranged differently to diamond. Graphite uses strong covalent bonds and weaker van der Waals forces.

In graphite, each carbon atom forms three single covalent bonds. A spare electron is left in the p orbital.

This arrangement produces a two dimensional layer of linked hexagons of carbon atoms.

The p orbitals with the "spare" electron merge above and below the plane of the carbon atoms in each layer. Electrons can move around the layer as they are delocalised which increases the strength of bonding as it is similar to metallic bonding only in 2 dimensions.

Graphite can conduct electricity due to the delocalised electrons. Can only conduct freely through the plane, not at right angles to them.

No covalent bonding between planes. Van der waals forces holds the layers together. Therefore layers can slide over each other.

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Bonding and Structure Summary

Ionic Bonding - Giant Structure, High melting point, conducts electricity in liquid and aqueous solution.

Covalent Bonding (giant macromolecular) - High melting point, does not conduct electricity.

Covalent Bonding (simple molecular) - Low melting point but may react in aqueous solution.

Metallic Bonding - Giant structure, high melting point, conducts electricity in solid form and liquid form. In aqueous solution, does not dissolve and may react.

Three types of intermolecular forces:

  • van der Waals - acts between all atoms
  • dipole-dipole forces - acts between molecules with a permanent dipole: X(delta+) - Y(delta -)
  • hydrogen bonding - acts between molecules formed when highly electronegative atoms (O, N or F) are covalently bonded with hydrogen atoms.
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The Shapes of Molecules and Ions

Electrons in molecules exist in pairs in volumes of space called orbitals. We can predict the shape of a simple covalent molecule (eg. one consisting of a central atom surrounded by a number of atoms) by using the ideas that:

  • each pair of electrons around an atom will repel all other electron pairs.
  • the pairs of electrons will therefore take up positions as far apart as possible to minimise repulsion.

This is called the electron pair repulsion theory. Electron pairs may be a shared pair or a lone pair.

The shape of a simple molecule depends on the number of pairs of electrons that surround the central atom. To work out the shape you first need to draw a dot-cross diagram to find the number of pairs of electrons.

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The Shapes of Molecules and Ions

If there are two pairs of electrons around the atom, it will be of a linear shape. The furthest the two pairs can get is 180 degrees apart. Example is Beryllium Chloride. (360 divided by two)

If there are three pairs of electrons around the central atom, they will be 120 degrees apart. The molecule is flat and the shape is trigonal planar. (360 divided by three). BF3 is an example of this.

If there are four pairs of electrons, the pairs spread far apart to form a tetrahedral shape. CH4 is an example of this. The bond angle is 109.5 degrees and it is a three dimensional shape. The sum of the angles can be more than 360 degrees. NH4- is also tetrahedral.

If there are five pairs of electrons, the shape is known as a trigonal bipyramid. It has three 120 degree angles on the right and two 90 degree angles on the left. Phosphorus pentachloride has this shape.

If there are six pairs of electrons, the shape adopted is octahedral with bond angles of 90 degrees. SF6 is an example of this.

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The Shapes of Molecules and Ions

Some molecules have unshared electrons (lone pairs). These are electrons which are not part of a covalent bond.

Lone pairs affect the shape of the molecule. Ammonia and water are examples where lone pairs affect the shape.

Ammonia - four pairs of electrons and one of the groups is a lone pair. With its four pairs of electrons, its shape is based on a tetrahedron but there are only three "arms". The shape is a triangular pyramid. 

You can also say that electron pairs form a tetrahedron but the bonds form a triangular pyramid.

The angles of a regular tetrahedron are all 109.5 degrees. However, lone pairs can affect these angles. In ammonia, the shared pairs of electrons are attracted towards the nitrogen nucleus and also the hydrogen nucleus. However, the lone pair is attracted only by the nitrogen nucleus and is pulled closer to it than the shared pairs. 

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The Shapes of Molecules and Ions

Repulsion between a lone pair and bonding pair is greater than that between two bonding pairs. This effect "squeezes" hydrogen atoms together, reducing the H-N-H angles. 

Approximate rule of thumb is 2 degrees per lone pair. So in ammonia, bond angles are approximately 107 degrees.

Water - Four pairs of electrons so the shape is of a tetrahedron. However, two of the "arms" are lone pairs that are not part of the bond. This results in a V-shaped molecule. Due to the two lone pairs, H-O-H angle reduces to 104.5 degrees.

Chlorine Tetrafluoride ion - Four bonding pairs and two lone pairs. One of the lone pairs has an electron which was donated to it, giving a -1 charge. 6 pairs of electrons (4 bonds, 2 lone pairs) - octahedron shape in which two arms are not part of the bond.

As lone pairs repel the most, they adopt a position furthest apart. This leaves a flat-square-shaped ion (square planar). The lone pairs are above and below the plane.

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The Shapes of Molecules and Ions

Bonding Pair - Bonding Pair    |

Lone Pair - Bonding Pair         |

Lone Pair - Lone Pair              |


Repulsion increases as you move down.

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Ta babes.

harshini vasik


thankyou ! very helpful

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