Physics Chapter 1

Chemical Bonds and the Structure of Substances

Electronic Structure of Noble Gases

Noble gases or Group 0 (or Group VIII) elements exist as atoms and are unreactive (inert) due to their stable electronic structures.

The outermost (valence) electron shell of these elements are fully-filled
e.g. Helium: 2
Neon: 2.8
Argon: 2.8.8

Atoms of other elements tend to lose, gain or share electrons in order to obtain the stable electronic structure of a noble gas. They do this by forming chemical bonds.

There are 3 main types of bonding –

Covalent bond: Very strong attraction between the nuclei of two atoms and the shared electrons. 2 characteristics:

Covalent bond is formed between atoms of the same element or atoms of different elements.

It is only formed between atoms of non-metals.

Examples of compounds with covalent bond: CH_{4} (methane), SO_{2} (sulphur dioxide), CO_{2} (carbon dioxide)

Ionic bond:Strong electrostatic attraction between positive and negative ions formed from the transfer of electrons from one atom to another.

Metals always form positive ions (via losing electrons eg. )

If non-metals form ions, they will produce negative ions (via gaining electrons).


Examples of compounds with ionic bond: NaCl (sodium chloride), Na_{2}O (sodium oxide), MgF_{2} (magnesium flouride)

Metallic bond:Strong attractive forces between the positive metal cations and the ‘sea of delocalised negative electrons’.

The number of electrons used to form chemical bonds is called the valency.

Covalent Bonds – Electron Sharing

Atoms share electrons in order to attain the electronic structure of a noble gas. When they share electrons, very strong covalent bonds are formed between the atoms.

There are 2 ways to represent a covalent bond: a dot-cross diagram or a line-bond diagram. In a dot-cross diagram, the electrons of one atom are represented by ‘x‘ while the electrons of another atom are represented by ‘●’. There are 2 electrons in each covalent bond with each atom contributing 1 electron to the bond.

A dative (coordinate) bond is a covalent bond where both electrons in the bond come from one of the atom and the other atom contributing no electrons.

Ionic Bonds – Electron Transfer

Positive ions (cations) are formed by removing electrons from atoms. In a cation, the number of electrons is less than the number of protons.

Negative anions (ions) are formed by adding electrons to atoms. In an anion, the number of electrons is more than the number of protons.

Metals form cations and non-metals form anions, so ionic compounds are formed when a metal reacts with a non-metal.

Ions are formed when electrons are transferred from a metal atom (which becomes a positive ion) to a non-metal atom (which becomes an a negative ion)

Chemical Bonds \& the Periodic Table

Elements in the same group have the same number of outer (valence) electrons, which means they have the same valency. Thus they form compounds (ions and molecules) of similar formulas with other atoms.

Form ions with similar formula e.g. Group I:Na^{+}, K^{+}

Group VII:Cl^{-}, Br^{-}

Form similar covalent molecules e.g.Group IV:CH_{4}, SiH_{4}

Form compounds with similar formula e.g.Group I with Group VII

NaCl, NaBr, KBr

(Simple) Molecular Substances

Most non-metallic elements (except noble gases) and covalent compounds have simple molecular structures consisting of small simple molecules e.g. water, nitrogen gas and glucose.

Between the two atoms within a simple molecule, there exist very strong intramolecular covalent bonds.

Between two molecules, there exist weak intermolecular forces of attraction.

Molecular structures have low melting, boiling points (usually below 200^{◦}C) [due to the weak intermolecular forces], generally do not conduct electricity in any state [due to the non-charged molecules present], usually insoluble (or low solubility) in water but soluble in organic solvents.

Ionic Compounds

Ionic compounds are made up of positive metal cations and negative non-metal anions.

These ions are held together by strong ionic bonds or electrostatic forces of attraction.

Large amounts of energy are required to break these bonds, so ionic compounds have high melting, boiling points.

Ionic compounds cannot conduct electricity in solid state. They only conduct electricity in the liquid (molten) state and aqueous (dissolved in water) state because the ions can move in these states.

Many of them dissolve in water but not in organic solvents.

Metals (Metallic Bonding)

In a metal, each atom gives up some electrons to become a positive ion (cation) [due to weak attraction of the nucleus for the outermost electrons].

Metallic bonding is the force of attraction between the positive metal cations and the ‘sea of negative electrons’

These small electrons are mobile, and go into the spaces in between the cations. The cations are big and cannot move (only vibrations).

Metals can conduct electricity in both solid \& aqueous state , because the electrons can move throughout the metal.

Metallic bonds are very strong, so metals usually have very high melting and boiling points.

The strength of the metallic bonds depends on the charge on the metal cation (or the number of valence electrons contributed to the sea of electrons per atom) and the ionic radius (size) of the cation.

Macromolecular (Giant Covalent) Structures

In a macromolecule, all the atoms are joined together by strong covalent bonds. There are no intermolecular forces of attraction within a macromolecule.

Macromolecular structures have high melting, boiling points [due to the strong covalent bonds], are insoluble in all solvents and do not conduct electricity (except for graphite) [due to the absence of charged particles]. Polymers like plastics are also macromolecules.

Examples of compound with macromolecular structures: Diamond, silicon, silicon dioxide (sand), graphite.

Diamond and graphite same atomic constituent but different physical properties

Diamond and graphite are allotropes of carbon.

In allotropes, atoms of the same element are arranged (bonded) in different ways.

Both of them have macromolecular structures. In diamond, each carbon atom is joined covalently to 4 other carbon atoms in a tetrahedral arrangement.

In graphite, each carbon atom is joined covalently to 3 other carbon atoms in flat, hexagonal layers. Between the layers are weak intermolecular (Van der Waals) attractive forces. The weak forces between layer account for the soft and slippery properties of graphite.

In addition, in graphite, one outermost electron of the carbon atom is not involved in covalent bonding. This delocalized electron is mobile and thus graphite can conduct electricity.

Silicon and silicon dioxide have macromolecular structure too. Silicon has a similar structure to diamond where each silicon atom is joined covalently to 4 other silicon atoms in a tetrahedral arrangement.

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