Trends in melting points of elements in Period 3

Melting points generally increase going from sodium to silicon, then decrease going to argon (with a “bump” at sulphur).

Explanation of the trends:

When a substance melts, some of the attractive forces holding the particles together are broken or loosened so that the particles can move freely around each other but are still close together. The stronger these forces are, the more energy is needed to overcome them and the higher the melting temperature.

Sodium, magnesium and aluminium

Sodium, magnesium and aluminium are all metals. They have metallic bonding, in which positive metal ions are attracted to delocalised electrons. Going from sodium to aluminium:
the charge on the metal ions increases from +1 to +3 (with magnesium at +2)
the number of delocalised electrons increases
so the strength of the metallic bonding increases and
the melting points increase.

Silicon

Silicon is a metalloid (an element with some of the properties of metals and some of the properties of non-metals). Silicon has giant covalent bonding. It has a giant lattice structure similar to that of diamond, in which each silicon atom is covalently-bonded to four other silicon atoms in a tetrahedral arrangement. This extends in three dimensions to form a giant molecule or macromolecule.
Silicon has a very high melting point because:
all the silicon atoms are held together by strong covalent bonds.
which need a very large amount of energy to be broken.

Phosphorus, sulphur, chlorine and argon

These are all non-metals, and they exist as small, separate molecules. Phosphorus, sulphur and chlorine exist as simple molecules, with strong covalent bonds between their atoms. Argon exists as separate atoms (it is monatomic).
Their melting points are very low because:
when these four substances melt or boil, it is the van der Waals’ forces between the molecules which are broken,
which are very weak bonds,
so little energy is needed to overcome them.
Sulphur has a higher melting point and boiling point than the other three because:
phosphorus exists as P4 molecules
sulphur exists as S8 molecules
chlorine exists as Cl2 molecules
argon exists individual Ar atoms
the strength of the van der Waals’ forces decreases as the size of the molecule decreases
so the melting points and boiling points decrease in the order S8 > P4 > Cl2 > Ar

What I've learnt this week:

1. We have to use Lewis Dot and Cross Diagram for covalent compounds.
   E.g: Ammonia:

• Lone pair of electrons: electrons are unshared
• Bond pair of electrons: electrons are shared.

2. Dative Covalent Bond:

• A type of covalent bond where the shared pair of electrons comes from one atom only.

3. Sigma bond, Pi bond

4. Shape of particular molecules

(continue next time)

Quick Recap: Chemical Bonding at 'O' Levels

1) Ionic bond: Formed between a metal and a non-metal.
It involves the transfer of electrons from a metal to a non-metal so that both can have noble gas configuration.

Ionic compounds have the following properties:

• High melting and boiling points, due to the strong electrostatic forces between oppositely charged ions, which require a lot of energy to overcome.

• Conducts electricity in molten (l) and aqueous (aq) states, but not in solid state. This is because in solid state, the ions are held in fixed positions in a lattice and cannot move around to conduct electricity. In molten and aqueous states, however, the ions are free to move around and carry charges.

• Soluble in water but not in organic solvents such as ethanol.

2) Covalent bonding: Sharing of valence electrons.

In covalent compounds, non-metals share their electrons so that each atom can have a noble gas configuration.

• Covalent bonds themselves are very strong. However, covalent compounds have low melting and boiling points. This is due to the weak intermolecular forces between the molecules, which require little energy to overcome.

Covalent compounds have the following properties:

• Covalent compounds cannot conduct electricity. They have no ions and no free electrons. Certain covalent compounds, however, can conduct electricity when dissolved in water. An example of such compounds would be acids. Hydrogen chloride (HCl) is a gas at room temperature, but dissolve it in water and you get hydrochloric acid. However, generally, covalent compounds cannot conduct electricity in whatever state.

• Covalent compounds don’t dissolve in water (with the exception of acids). They do dissolve in organic solvents.

3) Metallic Bonding: Bonding present WITHIN metal.
It consists of strong electrostatic forces between the positively charged metal ions and the delocalised electrons. The metallic structure consists of a lattice of positive ions in a sea of delocalised electrons.

• How a metal conducts electricity?

=>When a metal is used in an electrical circuit, the valence electrons entering one end of the metal causes a similar number of electrons to be displaced from the other end. The valence electrons move from the negative terminal to the positive terminal of the electric circuit.

The chemical bonding table Ms Jee gave was vey helpful.

Atomic Structure

• When forming cations, the 4s electrons will be lost first before the 3d
  electrons, even though 4s orbital is filled before the 3d orbital.

WHY?

=>The 3d orbital penetrates the electron density in the 4s orbital
considerably, thus shielding the 4s electrons from the attraction of the
nucleus and causing them to be easily lost.

=> The 3d subshell is closer to the nucleus. Once occupied by electrons,
it repels the 4s electrons even further from the nucleus and up to a higher
energy level.

Atomic Structure:

s subshells hold 2 electrons (or 1 orbital),
p subshells hold 6 electrons (or 3 orbitals),
d subshells hold 10 electrons (or 5 orbitals),
f subshells hold 14 electrons (or 7 orbitals).

The subshells are arranged in the following order:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, ...

• Atoms are ionised when they lose an electron. The energy required to remove the electron is known as the ionisation energy.
• As each electron is removed from an atom the ionisation energy required increases, so we call the energy required to remove the first electron the first ionisation energy,
• The energy required to remove the second electron the second ionisation energy .

First ionisation energy: the amount of energy needed to remove one mole of electrons from one mole of gaseous M atoms to form one mole of positively charged M+ ions.

E.g: M(g) => M+(g) + e-

Second ionisation energy: the amount of energy needed to remove one mole of electrons from one mole of gaseous ions to form one mole of gaseous M2+ ions.

E.g: M+(g) => M2+(g) +e-

• When electrons in the atom occupy the lowest energy levels, the atomis said to be in its ground state. Most atoms are in their ground state at room tempertature.
• When one or more electrons absorb enough energy to become promoted to a higher energy level, the atom is unstable and is said to be in an excited state.

Factors affecting Ionisation Energies:

• Magnitude of the positive nuclear charge.
• Shielding(Screening) effect of the Inner Electrons.

3 things I've learnt:

1)When a metallic element and a nonmetallic element combine, the nonmetallic atoms often pull one or more electrons far enough away from the metallic atoms to form ions. The positive cations and the negative anions then attract each other to form ionic bonds.

2)Metal‑nonmetal: Ionic compounds whose formula contains one symbol for a metal and one symbol for a nonmetal are called binary ionic compounds.

3)Metal‑polyatomic ion: Polyatomic ions can take the place of monatomic anions, so formulas that contain a symbol for a metallic element and the formula for a polyatomic ion represent ionic compounds.