intramolecular forces are ionic bonds,polar-covalent bonds,and non-polar covalent bonds. You can think

of these 3 types of bonds lying on a scale of unequal sharing of electrons. Ionic bonds

have the most unequal sharing, with one ion having a full negative charge, with one or

more extra electrons, and the other ion being fully positive since it has lost an electron

or two. Polar covalent bonds have more equal sharing of electrons, but one atom keeps the

electrons more towards their side of the bond than the other.

And non-polar covalent bonds

are between atoms that more or less equally share their electrons.

This video will talk

about the relative strengths of these 3 types of bonds.

Remember an ionic bond is formed from the electrostatic attraction between two oppositely

charged ions - an anion attracted to a cation.

Covalent bonds, on the other hand, involve

sharing of valence electrons. The shared electrons spend part of the time next to one of the

atoms in the bond, and part of the time next to the other atom.

In a polar covalent bond,

the more electronegative atom gets more of the electron’s time.

In a nonpolar covalent bond,

the electronegativities are very similar, and the electrons are shared equally.

Do we know which is stronger, ionic or covalent bonds? It seems a simple question, but there

is not one simple answer. One way we might determine the strength of a bond is by measuring

how much energy is required to disrupt it. NaCl, our example of an ionic bond, doesn’t

usually exist as a compound of just one sodium ion and one chloride ion (what we call a formula

unit). Sodium chloride typically exists as a lattice of many atoms. To disrupt that lattice

of NaCl into Na+ ions and Cl- ions requires 788 kJ/mol.

But we wouldn’t say that 788 kJ/mol

is the strength of ALL ionic bonds - that’s only the measured value for NaCl

in particular.

Rather, there is a range of lattice energies for different ionic compounds, which tells

us there are varying strengths of ionic bonds. This has to do with the charge of each ion,

which determines the strength of the attraction, and the radius of each ion, which influences

how closely the ions can be held together. In general, higher charge and smaller radii

lead to stronger bonds.

For instance, sodium iodide, NaI, made of Na+ and I-, has a lattice energy of 682 kJ/mol.

The ionic radius of iodide is larger than the ionic radius of chloride, and so NaI has

a slightly weaker bond than NaCl, because the ions aren’t held as closely together.

Magnesium oxide, MgO, on the other hand, is formed from Mg2+ and O2-. The higher charges

mean a stronger attraction, and MgO has a lattice energy of 3,795 kJ/mol.

Covalent bonds also vary in their strengths. Again, we can compare these bond strengths

by measuring the energy required to break the bonds. We call these “bond enthalpies,”

in the case of covalent bonds, and just like for lattice energies in ionic compounds, bond

enthalpies for covalent molecules are influenced by such factors as the distance the atoms

are held from each other and the differences in their electronegativities.

A single covalent bond between carbon atoms, for instance, has a bond enthalpy of 348 kJ/mol,

while a double bond between carbon atoms has a bond enthalpy of 614 kJ/mol.

Double bonds,

where atoms share 2 pairs of electrons, are shorter and stronger than single bonds. Triple

bonds are shorter and stronger still. A triple bond between two carbon atoms has a bond enthalpy

of 839 kJ/mol.

So you can see, there is a range of strengths, but Ionic and covalent bonds are both considered

strong bonds, the intramolecular forces that hold a compound together. In contrast, intermolecular

forces, such as van der Waals forces, which exist between different compounds, are examples

of bonds which are much weaker. We’ll talk about these weaker bonds in another video.