Melting points and boiling points
Melting points and boiling points depend on the strength of forces which exist between the particles which make up a substance.
Melting points and boiling points show periodic properties. This means that they vary in a regular way or pattern depending on their position in the Periodic Table.
Melting points and boiling points of Period 2 and Period 3 elements
The graphs above show the melting and boiling points for the period 2 and period 3 elements.
Melting points and boiling points of Group 1 metals
A trend can also been seen on descending a group.
The melting and boiling points of the Group 1 elements decrease on descending the group. This is due to a decrease in the forces of attraction between the atoms.
Atomic size
- On crossing a period in the Periodic Table the atomic size decreases.
- On descending a group the atomic size increases.
Crossing a period
Moving from one element to the next, electrons are being added to the same energy level and protons are being added to the nucleus.
The electrons in the outer energy level are therefore attracted more strongly and pulled closer to the nucleus.
This results in the atomic size decreasing.
Descending a group
Moving from one element to the next, the number of occupied energy levels is increasing.
This results in the atomic size increasing.
Ionisation energy
The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms.
Na(g) Na+(g) + e-
The equation above shows the first ionisation energy of sodium.
The second and subsequent ionisation energies refer to the energies required to remove further moles of electrons.
Descending a group
The first ionisation energy decreases on going down a group.
This is because the electron to be removed from the outer energy level is increasingly distant from the nucleus as a result of the atoms getting bigger.
Electron arrangement of Group 1 metals
The attraction of the nucleus for the electron becomes less and it becomes easier to pull it away.
Electrons in the inner energy levels also produce a screening effect. These inner electrons reduce the attraction of the nucleus for the outer electrons. The screening effect will increase as the number of inner energy levels increases.
Crossing a period
The first ionisation level generally increases on going across a period.
This is because on crossing a period, more protons are being added to the nuclei of the atoms. This results in an increase in nuclear charge.
The electrons in the outer energy levels will be more tightly held and more difficult to remove.
Electronegativity
Atoms of different elements have different attractions for bonding electrons.
Electronegativity is a measure of the attraction an atom involved in a bond has for the electrons of the bond.
An atom with a high electronegativity will tend to attract bonded electrons towards it.
An atom with a low electronegativity will have a very weak attraction for electrons.
Patterns in electronegativity
Electronegativity values:
- increase from left to right across a period in the Periodic Table
- decrease on descending a group
Be aware!
Electronegativity values can be found in the SQA data book (page 10 ).
Electronegativity values can be useful in predicting which type of bonding is most likely between two elements (see Bonding, structure and properties).
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Types of bonding
The different forms of bonding considered here are:
- metallic
- pure covalent
- polar covalent
- ionic
Metallic bonding
Metallic bonding is the type of bonding found in metal elements. This is the electrostatic force of attraction between positively charged ions and delocalised outer electrons.
Pure covalent bonding
Atoms in a covalent bond are held together by electrostatic forces of attraction between positively charged nuclei and negatively charged shared electrons.
Electronegativity (see Patterns in the Periodic Table) is the measure the attraction of an atom involved in a bond has for the electrons of the bond.
When two atoms bonded by a covalent bond have the same electronegativity the electrons will be equally shared. This is called pure covalent bonding. For example, the chlorine molecule, Cl2. Pure covalent bonding is only found in elements.
Polar covalent bonding
In most covalent compounds the bonding is polar covalent.
Covalent bonds with unequal electron sharing are called polar covalent bonds.
The atom with the higher electronegativity will have a greater share of the electrons. This will result in the atom having a slight negative charge compared to the other atom which will have a slight positive charge.
Ionic bonding
Ionic bonding is the electrostatic force of attraction between positively and negatively charged ions.
These ions have been produced as a result of a transfer of electrons between two atoms with a large difference in electronegativities.
Ionic bonding results from metals combining with non-metals. For example, sodium (a metal) combining with chlorine (a non-metal).
Be aware!
Covalent, polar covalent and ionic bonding can be considered as forming a spectrum. As the difference in electronegativity increases the bonds become more polar then ionic.
Intermolecular forces of attraction
Intermolecular forces of attraction are forces between different molecules.
The intermolecular forces of attraction considered here are:
- Van der Waals' forces
- permanent dipole-permanent dipole interactions
- hydrogen bonding
Van der Waals' forces
Van der Waals' forces are forces of attraction which exist between all atoms and molecules. Van der Waals' forces are much weaker than all other types of bonding.
They are only significant in atoms and molecules which have no other types of intermolecular forces of attraction, for example, discrete non-polar molecules and the Group 0 elements.
Van der Waals' forces are a result of electrostatic attraction between temporary dipoles and induced dipoles caused by movement of electrons in atoms and molecules.
The strength of Van der Waals' forces is related to the size of atoms and molecules. The bigger the atom or molecule the bigger the Van der Waals' force.
Permanent dipole-permanent dipole interactions
Permanent dipole-permanent dipole interactions are additional electrostatic forces of attraction between polar molecules.
Permanent dipole-permanent dipole interactions are stronger than Van der Waals' forces for molecules of equivalent size.
A molecule can be described as polar if it has a permanent dipole. A permanent dipole is due to a difference in electronegativity between the atoms involved in a covalent bond.
For example:
The spatial arrangement of polar covalent bonds can result in a molecule being polar.
If the molecule is symmetrical the molecule will be non-polar.
Hydrogen bonding
Hydrogen bonds are permanent dipole-permanent dipole interactions.
Bonds consisting of a hydrogen atom bonded to an atom of a strongly electronegative element such as fluorine, oxygen or nitrogen are highly polar.
Hydrogen bonds are electrostatic forces of attraction between molecules which contain these highly polar bonds.
Be aware!
A hydrogen bond is stronger than other forms of permanent dipole-permanent dipole interactions but weaker than a covalent bond.
Structures
The different types of structure considered here are:
- metallic
- covalent molecular
- covalent network
- ionic
- monatomic
Metallic structure
A metallic structure consists of a giant lattice of positively charged ions and delocalised outer electrons.
Covalent molecular structure
A covalent molecular structure consists of discrete molecules held together by weak intermolecular forces.
Covalent network structure
A covalent network structure consists of a giant lattice of covalently bonded atoms.
Ionic structure
An ionic structure consists of a giant lattice of oppositely charged ions.
Monatomic structure
A monatomic structure consists of discrete atoms held together by Van der Waals' forces.
The first 20 elements in the Periodic Table can be categorised according to bonding and structure.
Compounds can adopt one of three structures in the solid state:
- covalent molecular: eg carbon dioxide and methane
- covalent network: eg silicon dioxide and silicon carbide
- ionic: eg sodium chloride and potassium nitrate
Properties
The melting points, boiling points, solubility and hardness/softness of elements and compounds are related to their bonding and structures.
The melting points of polar substances are higher than the melting points of non-polar substances with similar sizes.
Ionic compounds and polar molecular compounds tend to be soluble in polar solvents such as water and insoluble in non-polar solvents such as carbon tetrachloride.
Non-polar molecular substances tend to be soluble in non-polar solvents and insoluble in polar solvents.
Boiling points, melting points, viscosity and miscibility in water are properties of substances which are affected by hydrogen bonding.
The boiling points of ammonia, water and hydrogen fluoride are higher than predicted due to the presence of hydrogen bonding between the molecules.
Hydrogen bonding between molecules in ice results in an expanded structure which causes the density of ice to be less than that of water at low temperature. This is why ice floats on water.