Structure and Bonding

1. Ionic Chemical Bonding

– Consider a specific case of ionic bonding between real elements, such as sodium and chlorine.
– A sodium atom contains 11 protons and has an electronic arrangement of 2.8.1.
– The arrangement differs from the nearest noble gas electronic structure, that of neon, 2.8 by the presence of one extra electron in the third energy level.
– On the other hand, a chlorine atom contains 17 protons and has the electronic arrangement of 2.8.7.
– It differs from the nearest noble gas electronic arrangement, that of argon, 2.8.8, by missing one electron in the third energy level.
– In order to attain the stable noble gas electron arrangement, a sodium atom would have to lose the electron in the outermost energy level.
– The chlorine atom would need to take one electron into its outer energy level to gain the noble gas structure.
– During ionic chemical bonding of sodium and chlorine atoms, the single electron from the outermost energy level of sodium atom is transferred to the outermost energy level of the chlorine as shown in Figure 1 below.

Figure 1: Ionic bonding in sodium chloride (NaCl)

– Sodium atom has 11 positive charges (protons) balanced by 11 negative charges (electrons).
– Sodium ion has only 10 electrons.
– Therefore, because the positive charges in the nucleus are unchanged, there is one excess positive charge on the sodium ion.
– Similarly, chloride ion has one negative charge in excess of the positive charges.
– A sodium ion is positively charged because of the one excess positive charge.
– Similarly chloride ion is negatively charged because of the one excess negative charge. These ions have opposite charges.
– Because of the attraction of the oppositely charged Na⁺ and Cl^– ions they attract and form a bond called ionic or electrovalent bond.
– This type of combination is called ionic bonding.

Figure 2: Sodium ion combine with chloride ion

Structure of Sodium Chloride

– The structure of sodium chloride contains numerous sodium and chloride ions in equal proportions.
– The electrical attraction electrostatic attraction resulting from their opposite charges constitutes the ionic bond.
– The ions arrange themselves into a rigid solid shape called a crystal.
– Each sodium ion is surrounded by six (6) equidistant chloride ions and vice versa.

Figure 3: Arrangement of sodium and chloride ions in sodium chloride crystal

– Sodium and chloride ions crystallise in a pattern (crystal lattice) forming a cube.
– In an end face of the cube a Na⁺ ion occupies the centre, with six Cl^– ions, spaced equally between them.
– The ions form a giant ionic structure.
– The attraction forces between the ions are strong and therefore the ions are not free to move but they vibrate within a given space.
– Consequently the melting point of sodium chloride is high. In solid form, it is a non-conductor of electricity.
– when sodium chloride is melted sodium ions and chloride ions separate and thus their forces of attraction are greatly reduced.
– When an electric current is applied the ions in molten sodium chloride are free to move thereby conducting electricity.
– The positive ions formed as a result of loss of one or more electrons are called cations, and their positive charges are equal to the number of electrons lost.
– Likewise, the negative ions formed as a result of gain of one or more electrons are called anions and their negative charges are equal to the number of electrons gained.
– The number of electrons lost from, or added to, the outermost energy level of the atom of an element during ionic bonding is equal to the combining power (valency) of that element.
– Only the outermost energy level electrons are involved in ionic bonding.
– The number of ions involved must balance the valency requirements of elements as shown in the following examples.

(i) Sodium sulphide

Figure 4: Ionic bond in sodium sulphide (Na₂S)

– Therefore the formula of sodium sulphide is Na₂S.
– The valency electrons (outermost energy level electrons) from the two sodium atoms are transferred to the outermost energy level of sulphur as shown in Figure 4 above.

(ii) Magnesium oxide

Figure 5: Ionic bond in magnesium oxide (MgO)

– Therefore the formula of magnesium oxide is MgO.
– The two valency electrons from one magnesium atom are transferred to one oxygen atom.

Properties of Ionic Compounds

Covalent Chemical Bonding

– Sharing a pair of electrons forms a covalent bond.

Examples

– A hydrogen molecule, H₂, has two hydrogen atoms linked by a covalent bond.
– A hydrogen atom has 1 electron in its only energy level. It is unstable.
– Therefore two hydrogen atoms combine by each contributing an electron each.
– Then they share the electron pair equally. The shared pair revolves around both atoms.
– In effect they have the stable electron duplet arrangement i.e. first energy level with 2 electrons like helium see Figure 6.

Figure 6: Hydrogen molecule

– The shared pair is the covalent bond. It is attracted by the proton of each H atom.
– Sometimes the shared pair is represented by a short line (–) H – H.
– A molecule of chlorine, Cl₂, contains two chlorine atoms linked by a covalent bond.
– These two chlorine atoms have each an electronic arrangement of 2.8.7 with 17 protons.
– If no other element is available from which electrons may be obtained to make these two atoms have a noble gas electronic structure such as argon (2.8.8), a shared pair of electrons is formed.
– Each chlorine atom contributes ,one electron to the shared pair.
– This idea of sharing can be shown in a diagram as in Figure 7.

Figure 7: Chlorine molecule

– In the chlorine molecule the stable electron octet arrangement for each chlorine atom is achieved.
– The nucleus of each atom attracts the shared pair strongly.
– The two atoms will remain joined together by each atom attracting the shared pair.
– Some atoms can share more than one pair of electrons.
– The covalency of an atom is the number of electron pairs which it shares.
– Atoms of different elements can also form covalent bonds like in carbon(IV) oxide. See Figure 8 below.

Figure 8: Illustration of various molecules

Properties of Covalent Compounds

Co-ordinate (Dative) Bonding

– A co-ordinate bond also called a dative bond is a covalent bond (a shared pair of electrons) in which both electrons come from only one atom.
– One of the atoms possesses a lone pair of electrons, i.e. a pair of electrons not bonded to any atom.
– This lone pair is shared with an atom or an ion that needs them to build up, or complete an electron octet or duplet and therefore attain stability.
– For example ammonia molecule which possesses such a lone pair of electrons is shown in Figure 9.
– This lone pair can be donated to hydrogen ion (H⁺) from an acid to form the ammonium ion, NH⁺₄.
– The nitrogen atom is said to be a “donor” and the hydrogen ion is said to be an acceptor.

Figure 9: Co-ordinate bond of ammonium ion

– The representation of a co-ordinate bond is →.
– The arrow points from the donor atom to the acceptor atom as shown above.
– The hydrogen ion contributes the charge on the ion formed.

Structures of Covalent Compounds

(i) Giant atomic structures

(a) Diamond

– In diamond, the carbon atoms are bonded to each other by covalent bonds in interlacing tetrahedrons throughout the structure.
– This produces a crystal which is a giant atomic structure.
– The covalent bonds between the atoms are very strong making diamond very hard.
– All four-valence electrons per atom are involved in covalent bond formation with four adjacent carbon atoms.
– The carbon atoms in diamond are close to each other and therefore diamond is dense.
– It has no free electrons delocalised electrons and therefore it does not conduct electricity.

Figure 10: Diamond structure

(b) Graphite

– In graphite, the carbon atoms are covalently bonded to each other in hexagonal rings arranged in parallel planes, one on top of the other, with the layers joined by weak van der Waals forces.
– The layers are able to slip over each other and this makes graphite soft and slippery (See Figure 11 below).
– The layers form a giant atomic structure.

Figure 11: Graphite structure

– In the parallel atomic layers (plates) of carbon atoms in graphite, only three valency electrons per atom are used in bond formation.
– Therefore, some electrons in graphite are free delocalised and allow it to conduct electricity.

(c) Silicon (IV) oxide

– Silicon is another good example involving atom to atom covalent bonding to form a giant atomic structure.
– Silicon(IV) oxide is insoluble in water and does not conduct electricity.
– The bonds between the oxygen – silicon atoms are strong and therefore silicon(IV) oxide has a high melting point.

Figure 12: Silicon(IV) oxide structure

(ii) Simple molecular structures

– Substances such as iodine, hydrogen, chlorine, nitrogen and carbon(IV) oxide exist as molecules.
– The atoms in the molecule are held by strong covalent bonds.
– However, the molecules are held together by weak van der Waals forces.
– Therefore when heated they easily vaporise or melt.
– An example is iodine structure.

Figure 13: The Structure of iodine

– On heating iodine, it sublimes because the weak van der Waals forces are easily broken.
Note: Molecules have discrete units with a definite number of atoms which are covalently bonded.
– Molecular structures (involving atoms or molecules) have very large indefinite number of atoms in them.

Figure 14: Molecules held together by van der Waals forces

Metallic Bonding

– The outermost energy level electrons in metals are relatively few.
– When the atoms of metals are closely packed, each metal atom loses its outer electron(s) which form a sea of free electrons (delocalised mobile electrons).
– The resulting metal positive ions are embedded in the sea of electrons.
– There is attraction between the ions and electrons.
– This kind of electrostatic attraction between the positive metal ions and the delocalised electrons form the metallic bond.
– The ions arrange themselves into a giant metallic structure.

Figure 15: Metallic bond

– These delocalised electrons can move on application of an electric current or heat.
– This explains why metals are good conductors of electricity and heat.
– Metallic bonding is very strong in some metals like copper, iron, but weak in others like sodium and potassium, which can be cut with a knife.
– Moving across a period (e.g. period 3) of the periodic table, the number of valency delocalised electrons increases and therefore the strength of the metallic bonds increases.
– Hence the melting points and boiling points of metals increase across the period.
– The thermal and electrical conductivity also increases across the period because of the increase in the number of valency electrons (delocalised electrons).

– The melting point increase from sodium to aluminium. The change in melting point from magnesium to aluminium is not very big.
– This is probably because not all the three electrons in aluminium are involved in metallic bonding.
– As the strength of metallic bond increases across the period, the pull of positive ions towards each other increases thus also increasing the density of the metals.
– It might be expected that increase in temperature would speed the movement of free electrons, with a consequent increase in electrical conductivity.
– In general, however, the electrical conductivity of metals decreases with increase in temperature.
– This is because increasing temperature produces increased thermal vibration within the metal structure.
– This upsets the regularity within the crystal and interferes with the ease of movements of electrons within the crystal.
– It is just like comparing soldiers matching on parade and others matching through a city crowd.

Hydrogen Bonding

– For example in a water molecule(H₂O), oxygen atom attracts electrons more than hydrogen.
– Therefore the water molecule is represented like this: H^δ+ – O^δ– – H^δ+ The water molecules can combine together using the polar ends.
– The attraction between the polar end of the hydrogen (δ+) and the polar end of the oxygen (δ -) is called hydrogen bond.

– Therefore water has hydrogen bonds between molecules.
– Had it not been the presence of hydrogen bonds, water would be in gaseous state at room temperature and pressure.

Types of Bonds Across Period Three

(i) Oxides of elements in period three
– Oxides of sodium, magnesium and aluminium form giant ionic structures.
– They have therefore high melting points. Sodium oxide is soluble in water, while magnesium oxide is only slightly soluble.
– When these two oxides dissolve they form basic solutions. They are therefore basic oxides.
– Sodium oxide reacts explosively with dilute mineral acids while magnesium oxide reacts at a reasonable rate to form salt and water only.
– Aluminium oxide is insoluble in water. It however reacts with both dilute mineral acid explosively and dilute alkalis to form salts.
– This oxide is said to be amphoteric, that is, it has both acidic and basic properties.
– Silicon(IV) oxide forms giant atomic structure, while phosphorus(V) oxide, sulphur(IV) oxide, chlorine(VII) oxide (Cl₂O₇) form molecular structures.
– The oxides of phosphorus and sulphur react with water to form phosphoric(V) acid and sulphurous acid respectively.
– They are therefore acidic oxides. Silicon(IV) oxide does not react with water or dilute alkali; but it reacts with concentrated alkalis to form salt and water only.
– Therefore it is regarded as an acidic oxide.
– The oxides of silicon, phosphorus, sulphur and chlorine have no effect on acids.

Chloride of Elements in Period Three

– Chlorides of sodium and magnesium form giant ionic structures which have high melting and boiling points.
– These two chlorides are soluble in water.
– The chlorides of aluminium, silicon, phosphorus and sulphur are simple molecular structures.
– These four chlorides react with water with evolution of heat to form acidic solutions and also give off fumes of hydrogen chloride gas.
– Some of the hydrogen chloride gas dissolves in water to form hydrochloric acid.
– This type of reaction is called hydrolysis and involves all molecular chlorides.

Applications of Diamond, Graphite, and Aluminium

Diamond

(i) Jewelry, it has a shiny lustre when polished.
(ii) Glass cutter and drilling because it is very hard.

Graphite

(i) As a lubricant, because it is soft.
(ii) Reinforcement of metals and broken bones.
(iii) As electrodes because it conducts electricity.

Aluminium

(i) Overhead electrical cables.
(ii) In making of cooking utensils.
(iii) Wrapping material.
(iv) Silvering of mirrors and different types of reflectors.
(v) Aluminium paint.

All these applications take advantage of the physical and chemical properties of these three elements. These properties are determined by the type of bonding between the atoms of the elements. For example, let us consider aluminium:
(i) It is a conductor of electricity. It is lighter than copper.
(ii) It is a good conductor of heat. It resists corrosion and has a low density.
(iii) It is soft and malleable.
(iv) It does not tarnish and is a very good reflector of light.
(v) It is a light metal.
(vi) The paint of aluminium is normally applied on iron or steel to prevent corrosion or rusting.