Bonds

Ionic

This occurs between oppositely charged ions (atoms or molecules which have gained or lost electrons). An example of an ionically bonded compound is NaCl.

The sodium exists as a positive ion, Na⁺. It loses the electron in its 3s¹ sub-shell, giving it a stable electron configuration of 1s² 2s² 2p⁶, or [Ne]. This electron is donated into the 3p sub-energy level of the chlorine atom, which then forms a chloride ion, Cl^-. This also now has a stable electron configuration of 1s² 2s² 2p⁶ 3s² 3p⁶ or [Ar].

The force of attraction (called the "electrostatic force of attraction") between these two ions, caused by their opposite charges, is the ionic bond.

The ions arrange themselves into "giant ionic lattices", which are regular, repeating structures. In NaCl, every Na⁺ ion is surrounded by 6 Cl^- ions, and every Cl^- ion is surrounded by 6 Na⁺ ions. Ionic compounds are usually white and crystalline in appearance. They have high melting and boiling points, as a lot of energy is required to overcome the electrostatic forces of attraction. They can, however, be disrupted by polar solvents, such as water. Ionic compounds are often, therefore, soluble in water. Ionic compounds are brittle - they shatter when forced to change shape. Ionic compounds do not conduct electricity when solid because there is no way for electrical charge to travel through the material. However if molten, or in solution, the ions can move and carry an electrical current.

Covalent

The term covalent bond is used to describe the bonds in compounds that result from the sharing of one or more pairs of electrons. This is usually indicated as (for example) H-F for a single pair of electrons, O=O for double bonds (two pairs of electrons) and N≡N for a triple bond. The electron pair creates a 'bond' between the two atoms because it attracts the nucleus of each atom and therefore resists the separation of the two atoms.

A co-ordinate bond (more correctly called a dative covalent bond) is a covalent bond in which both electrons come from the same atom. Coordinate bonds are shown using an arrow from the atom that doated the two electrons. For example, :NH₃ has a lone pair of electrons it can use to form a bond with BH₃, where the boron atom only has 6 outer electrons. The bond between N and B is shown as H₃B←NH₃.

Notice how, where possible, we alter the chemical formula to help show which atom bonds to which. We could write BH₃←NH₃ but it does not help to show that it is the N bonding to the B.

Another example: Two :NH₃ molecules will form coordinate bonds with an Ag⁺ ion. The ion that forms has the structure [H₃N→Ag←NH₃]⁺.

Metallic

Positive ions are arranged in a lattice with a sea of delocalised (free) electrons. The force of attraction between the delocalised electrons and positively charged ions holds the structure together. See the next chapter for full details

Intermolecular Forces

The bonds which act between a molecules are called intermolecular forces. The word intermolecular means, "between molecules". There are three kinds of intermolecular forces.

Van Der Waals Forces

Van der Waals forces are forces which attract two molecules together without actually forming bonds. This definition excludes ionic bonding and hydrogen bonding. Some sources, incorrectly, also exclude dipole-dipole forces or include hydrogen bonding.

Induced Dipole-Induced Dipole Forces

Also known as ID-ID forces, Instantaneous Dipole-Induced Dipole forces, Temporary Dipole-Induced Dipole forces, London forces or Dispersion forces. These are the weakest of the intermolecular forces, but they are always present, and even pull atoms of noble gases together.

ID-ID forces are stronger if the molecules have more electrons, or if they have larger surface areas. For example, the boiling points of propane, butane and methylpropane are -42 °C, -0.5 °C and -12 °C respectively. Propane has the lowest boiling point; It is the smallest molecule, with the fewest electrons and the smallest surface area. Butane and methylpropane are isomers, with the same number of electrons (34) in each molecule. Methylpropane has a smaller surface area than butane because the methylpropane molecule is more branched - it is closer to being spherical than the long butane molecule.

Propane boils at -42 °C
Butane boils at -0.5 °C
MethylPropane boils at -12 °C

Dipole-Dipole Forces

If two molecules are both dipoles, they will attract one another; The positive end of one dipole will attract the negative end of the other dipole, and vice-versa. These are stronger forces than ID-ID forces.

Dipole-Dipole forces are occasionally called Keesom forces.

Two propanone ("acetone") molecules attract one another if their dipoles (δ+ to δ- partial charges) are aligned.

Hydrogen Bonding

This is the strongest type of intermolecular force. Hydrogen bonds only occur between a hydrogen atom and a lone pair of electrons on a nitrogen (N), oxygen (O) or fluorine (F) atom. The H atom must have a strong partial charge (δ+), and this will be because the H atom is bonded to another N, O or F atom. The electronegative atom (N, O or F) pulls electrons away from the hydrogen so that, on the opposite side to the bond, the hydrogen appears almost like an unshielded proton.

Hydrogen bonds have a particular geometry. The N, O or F lone pair will have a specific angle relative to the other electron pairs, according to VSEPR theory. The X-H···X angle has to be close to 180°

Shapes of Molecules - VSEPR Theory

Wikipedia has related information at molecular geometry.

Valence shell electron pair repulsion theory (VSEPR) is used to predict the shape of a molecule. "Valence shell electrons" are simply the electrons in the outer shell of an atom. Electrons are usually in pairs, and they repel one another. The shape of the molecule is a natural consequence of electrostatic repulsion between every electron pair.

You can use the so-called AXE method to calculate the shape of a molecule. "A" represents which atom's geometry we want to analyse. "X" represents anything bonded to atom A, either atoms or groups of atoms, whether bonded with a single, double or triple bond. Lone pairs are labelled E. A molecule with three lone pairs and two atoms/groups bonded to it would be denoted AX₂E₃. The table below shows how X and E and molecular shape are related.

AXE label	X (substituents)	E (lone pairs)	Shape	Examples
AX₁E₀	1	0	Linear	H₂
AX₂E₀	2	0	Linear	BeCl₂ HgCl₂ CO₂
AX₁E₁	1	1	Linear	CN⁻
AX₃E₀	3	0	Trigonal planar	BF₃ CO₃²⁻ NO₃⁻ SO₃
AX₂E₁	2	1	Bent	NO₂⁻ SO₂ O₃
AX₁E₂	1	2	Linear	O₂
AX₄E₀	4	0	Tetrahedral	CH₄ NH₄⁺ PO₄³⁻ SO₄²⁻ ClO₄⁻
AX₃E₁	3	1	Trigonal pyramidal	NH₃ PCl₃
AX₂E₂	2	2	Bent	H₂O H₂S OF₂\|-
AX₁E₃	1	3	Linear	HCl
AX₅E₀	5	0	Trigonal Bipyramidal	PCl₅
AX₄E₁	4	1	Seesaw	SF₄
AX₃E₂	3	2	T-shaped	ClF₃ BrF₃
AX₂E₃	2	3	Linear	XeF₂ I₃⁻
AX₆E₀	6	0	Octahedral	SF₆
AX₅E₁	5	1	Square pyramidal	ClF₅ BrF₅
AX₄E₂	4	2	Square Planar	XeF₄

Methane molecule

There are multiple ways to show the structure of a molecule like methane. Dot-cross diagrams show each pair of electrons in the outer shell (the "valence electrons"). The 3D shape depends on the number of independent valence electron pairs.

dot-cross diagram
displayed structure
ball-and-stick model
spacefilling model
representation with 'wedge' and 'hash' bonds
connecting the H-atoms shows the tetrahedral structure

Ammonia molecule

dot-cross diagram
ball-and-stick model with lone pair displayed
ball-and-stick model without lone pair displayed
structural drawing with lone pair displayed

Water molecule

dot-cross diagram
ball-and-stick model with lone pairs displayed
ball-and-stick model without lone pairs displayed
structural drawing with lone pairs displayed

Tetrahedra

You have probably come across tetrahedra before in maths, although you most likely called them triangle-based pyramids. Tetrahedra have four vertices (corners), four faces and six edges. Each face is an equilateral triangle.

The tetrahedron is one of the most important shapes in chemistry because a very great many molecules contain them. Tetrahedral molecules don't actually contain little pyramids. What they do contain is a central atom bonded to four other atoms. The four atoms surrounding the central atom occupy positions that you can imagine as the vertices of a tetrahedron.

In the image gallery below, the central atom is coloured magenta and the surrounding atoms are coloured white.

File:Tetrahedron-2-3D-balls.png
File:Tetrahedron-4-3D-balls.png
File:Methane-2D-small.png
File:Methane-2D-flat-small.png

The angle between any two bonds in a tetrahedral molecule is approximately 109.5°. The tetrahedral angle can be calculated as accurately as required because it is equal to cos⁻¹(–⅓).

Octahedra

You may or may not have met an octahedron before. Octahedra have six vertices (corners), eight faces and twelve edges. Each face is an equilateral triangle.

Octahedra are very important in chemistry because many complex ions formed by metals are octahedral.

an octahedral molecule
to see the octahedron, connect the surrounding atoms with lines
the lines form the edges of the octahedron
octahedral angle = 90° exactly
chemists represent octahedra using hashed and wedged bonds
this is how chemists represent sulfur hexafluoride, SF₆
a ball-and-stick model of SF₆

Common molecular geometries

linear
X-A-X angle: 180°
bent
X-A-X angle: 104.5° (in H₂O)
trigonal planar
X-A-X angle: 120°
trigonal pyramidal
X-A-X angle: 107° (in NH₃)
square planar
X-A-X angle: 90°
tetrahedral
X-A-X angle: 109.5°
trigonal bipyramidal
X-A-X angles: 90° and 120°
octahedral
X-A-X angle: 90°

Further examples

Example molecules

linear: CO₂
bent: H₂S
trigonal planar: BI₃
pyramidal: NH₃
square planar: XeF₄
tetrahedral: CH₄
trigonal bipyramidal: PCl₅
octahedral: SF₆

The shape of a molecule can affect its polarity. CO₂, BI₃, XeF₄, PCl₅ and SF₆ all have polar bonds. However, the geometry of the bonds is symmetrical and the polarity of the bonds in each molecule will cancel out, leaving a non-polar molecule.

A polar molecule must have polar bonds and an asymmetrical structure. CF₄ and CH₂Cl₂ are similar molecules. The C-F bonds are more polar than the C-Cl bonds (and C-H bonds are hardly polar at all). However, CF₄ is symmetrical and therefore it is non-polar. CF₄ boils at -127.8 °C and CH₂Cl₂ boils at 39.6 °C. The lighter, but asymmetrical, molecule CH₂F₂ boils at -51.7 °C.

Example ions

linear:
hydroxide, OH⁻
trigonal planar:
carbonate, CO₃²⁻
trigonal planar:
nitrate, NO₃⁻
trigonal pyramidal:
hydronium, H₃O⁺
tetrahedral:
thiosulfate, S₂O₃²⁻
tetrahedral:
sulfate, SO₄²⁻
tetrahedral:
phosphate, PO₄³⁻
tetrahedral:
ammonium, NH₄⁺

Answering Questions

Many questions on this topic will ask about two or more chemicals and give you information such as:

Chemical A has a higher boiling point than chemical B
Chemical A has a higher melting point than chemical B
Chemical A is less volatile than chemical B
Chemical A is a solid and chemical B is a gas
Chemical A is a solid and chemical B is a liquid
Chemical A is a liquid and chemical B is a gas

If the chemicals are small molecules, in every case, you can immediately say that Chemical A has stronger intermolecular bonds than Chemical B.

If the chemicals are ionic, metallic or giant covalent structures, you can immediately say that Chemical A has strong bonds than Chemical B.

Now, you need to identify the bonds. If the chemicals have different types of bond, you have an answer: For example, "Chemical A has hydrogen bonds which are stronger than the dipole-dipole forces in chemical B".

If the chemicals have the same bonds, then you need to dig deeper, why are these bonds stronger in chemical A? Dipole-dipole bonds depend on the strength of the dipole i.e. the differences in electronegativity. Induced dipole-induced dipole forces are stronger if the molecules have larger surfaces and/or more electrons. Metallic bonds are stronger if the metals release more electrons and/or form smaller ions (compare group 1, 2 and 3 metals). Ionic bonds are stronger if the ionic charges are large and/or the ions are small.