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Chemical Bonding

  • The noble gases which have fully filled shells do not undergo chemical combination.
  • Other atoms undergo chemical combination in order to have fully filled shells and becomes stable like the noble gases nearest to them in the Periodic Table
  • The two main methods of chemical combination are Ionic (Electrovalent) Bonding, and Covalent Combination
  • Coordinate (Dative) Bonding is a type of covalent combination
Ionic Bonding
  • Ionic bonding always takes place between a metal atom and a non-metal atom
  • It involves a complete transfer of electrons from the outermost shell of a metal atom to the outermost shell of a non-metal atom
  • The number of electrons lost by the metal atom is always equal to the valency of the metal
  • The number of electrons accepted by the non-metal atom is equal to the valency of the non-metal atom
  • Ionic bonding results in the formation of discrete ions – the metal atoms becomes positively charged while the non-metal atom becomes negatively charged
  • Each positive ion carries a charge equal to its valency and each non-metal atom also carries a charge equal to its valency
  • Ionic bonding between magnesium and oxygen atoms may be illustrated as follows:

    Photo Credit: GCSE Science

    • Magnesium now has the electronic configuration of the noble gas neon while the oxygen ion also has the electronic structure of neon.
    • The electrostatic attraction between the opposite charged ions constitutes the ionic (electrovalent) bond.
  • In general, two kinds of atoms will undergo ionic bonding if the difference between their electronegativity values is greater than 1.7
Covalent Bonding
  • Covalent bonding always occurs between non-metal atoms
  • It involves a sharing of electrons contributed in equal numbers by the participating non-metal atoms which end up having fully filled shells like the noble gases nearest to them
  • The number of electrons contributed by a non-metal atom for sharing is equal to the valency of the non-metal atom
  • Each shared pair of electrons is equivalent to a single covalent bond which is usually represented as a single stroke
  • Two shared pairs of electrons are equivalent to two single covalent bonds or one double covalent bond
  • Three shared pairs of electrons make up three single covalent bonds or a triple bond

Examples of Covalent Combination

  • Formation of hydrogen molecule, H2

    Photo Credit: Dummies.com

    • The valency of H is 1. Therefore each H atom contributes 1 electron for sharing
    • In the image above, (b) is the dot notation way of showing the bond, and (c) shows the type of covalent bond
  • Formation of Nitrogen molecule, N2

    Photo Credit: (a) GSCE Science (b) Dummies.com

    • Each nitrogen atom exhibits a valency of 3 and so each atom contributes 3 electrons towards the formation of the triple bond between them
  • Formation of ammonia molecule, NH3

    Photo Credit: Zigya.com

    • The nitrogen atom contributes 3 electrons but each hydrogen atom contributes only 1 electron to account for the formation of 3 single covalent bonds in the ammonia molecule
  • A lone pair of electrons is a pair of electrons in the outermost shell of an atom which is not used for covalent bond formation
  • There are two lone pairs of electrons on the oxygen atom in a molecule of water

Coordinate Covalent Bonding

  • Coordinate covalent bonding always involves the sharing of lone pair of electrons donated by an atom in a molecule when it combines with another atom or ion that does not have a fully filled outermost shell. Some examples include:
Properties of Ionic Compounds
  • They conduct electricity in the molten state or in aqueous solution when their oppositely charged ions are mobile
  • They dissolve readily in water (polar solvent) but not in non-polar solvents
  • They are generally solids of high melting points because their oppositely charged ions are strongly held together by electrostatic forces
  • They form giant lattices
Properties of Covalent Compounds
  • They do not conduct electricity in the molten state because they do not contain ions
  • They do not dissolve in polar solvents like water, but usually dissolve in non-polar solvents
  • They have low melting and boiling points because of the weak inter-molecular forces of attraction between their molecules
  • They are usually liquid, gas, or low melting point solids
Metallic Bonding
  • The atoms of every metal are held together by metallic bonds
  • Every atom in a metal contributes its valence electrons to produce a ‘pool’ or ‘sea’ of electrons
  • This sea of electrons move randomly through the lattice of the residual positive ions
  • The free or delocalized electrons act like a glue between the residual ions and prevent them from repelling one another
  • The forces of attraction between the positively charged metal ions and the sea of electrons moving through the lattice constitute the metallic bond
  • The strength of the metallic bond increases with the number of valence electrons in the metal
  • Metallic bonding is responsible for the hardness of metals and their high melting points. However, both potassium and sodium are so soft that they can be cut with a knife
  • The delocalized nature or the mobility of the sea of electrons in metals accounts for their malleability and ductility as well as their ability to conduct heat and electricity

Photo Credit: Byjus.com

Inter-molecular Forces
  • Ionic bonds, covalent bonds, and metallic bonds are inter-atomic forces because they operate between atoms
  • Inter-molecular forces are weak forces which operate between molecules
  • These inter-molecular forces include Van dee Waal forces, Dipole-dipole attractions, and Hydrogen bond

Van der Waal Force of Attraction

  • This is the weakest inter-molecular force which operates mainly in non-polar molecules
  • It is actually the attraction between induced dipoles which develop when non-polar molecules experience instantaneous polarization.
  • Van der waal force of attraction increases with molar mass or the electron density within a system
  • Van der waal force of attraction is the main inter-molecular force of attraction in hydrocarbon molecules
  • As the molar mass increases in the alkane homologous series, Van der wall force increases and so do the physical properties, such as boiling point, density, and viscosity
  • Van der waal force of attraction explains why fluorine and chlorine are gases while bromine and iodine, with much larger molar masses, are liquid, and solid respectively.

Hydrogen Bond

  • Hydrogen bond is the strongest inter-molecular force of attraction
  • It is a special dipole-dipole attraction found only in molecules that have hydrogen atoms directly attached to strongly electronegative atoms such as nitrogen, oxygen, and fluorine.
  • In such compounds, such as HF, H2O, the hydrogen atom carries a partial positive charge which serves as a point of attachment to another electronegative atom
  • The strongest hydrogen bond is found in hydrogen fluoride because fluorine has the highest electronegativity
  • Hydrogen bond has a pronounced effect on the boiling point and the solubility in water of any substance in which it occurs
    • Water, molar mass of 18, has a boiling point of 100 °C because its molecules are joined together by hydrogen bonds
    • Hydrogen sulfide, molar mass 34, has a boiling point of -61 °C because its molecules are held together by dipole-dipole attractions which are weaker than hydrogen bonds
  • The presence of hydrogen bonds in water accounts for the high surface tension in water. It also accounts for the lower density (or larger volume) of ice which floats on water
  • Some covalent solutes or organic compounds, such as glucose, and ethanol, which are polar, dissolve readily in water because they can engage in hydrogen bonding with water molecules
  • Hydrogen bonds are always denoted by broken lines joining molecules of compounds exhibiting hydrogen bonding

Photo Credit: Socratic.com

Hybridization and Shapes of Molecules
  • Hybridization is the blending up of atomic orbitals to produce equivalent hybrid orbitals which are oriented in space at definite angles
  • The common types of hybridization include SP3, SP2, SP

SP3 Hybridization

  • This type of hybridization is characteristic of alkanes and may therefore be illustrated by the formation of methane.
  • In its ground state, the electronic structure of carbon is given by 6C: 1S2 2S2 2Px1, 2Py1, 2Pz0
  • In its excited state, after absorbing energy, carbon promotes one of its 2s electrons into the 2Porbital. Thus, its excited state electronic structure is given by 6C: 1S2 2S1 2Px1, 2Py1, 2Pz1
  • For SP3 hybridization, carbon blends up its 2s orbital with the three 2p orbitals to form four SP3 hybrid orbitals which are spread apart in space at 109°28′ apart in order to minimize repulsion
  • The spatial arrangement of these four SP3 hybrid orbitals gives rise to a tetrahedral shape that is characteristic of every alkane
  • To form methane, CH4, each of the four hydrogen atoms uses its 1s orbital to overlap with an SP3 hybrid orbital
  • In general, the greater the extent of overlap of atomic orbitals, the larger the amount of heat evolved and the stronger the bond formed
  • The extent of the co-linear overlap between each SP3 hybrid orbital and the 1s orbital of hydrogen in methane is the same
  • Consequently, the amount of heat liberated in each case is the same and hence each of the four C – H bonds has the same bond energy
  • Any bond formed by co-linear overlap of atomic orbitals is called a sigma bond (σ – bond)
  • Without hybridization, all four C – H bonds in methane would not have the same bond energy

Photo Credit: Nigerian Scholars

SP2 Hybridization

  • This type of hybridization is characteristic of alkenes such as ethene, C2H4
  • In its excited state, each of the two carbon atoms in ethene blends up its 2s orbital with only two (the 2Px and 2Py) of its three 2P orbitals to form three SP2 hybrid orbitals which are spaced at 120° apart in space
  • Each carbon atom uses two of its SP2 hybrid orbitals to overlap with two hydrogen atoms
  • These overlaps result in the formation of two sigma bonds between each carbon atom and two hydrogen atoms
  • An overlap of the third SP2 hybrid orbital between the two carbon atoms results in the formation of a sigma bond between them
  • The two carbon atoms form another bond called a pi bond between them through the overlap of the unused 2Pz orbitals on the two carbon atoms
  • The spatial arrangement of the SP2 hybrid orbitals causes every alkene to have a bond angle of 120° and a trigonal planar shape
  • The double bond between the two carbon atoms in ethene is made up of one sigma bond and one pi bond
  • The pi bond is weaker than the sigma bond and it is the point of attack (as it is easily broken) when alkenes undergo addition reactions
  • This is why alkenes are more reactive than alkanes

Photo Credit: Nigerian Scholars

SP Hybridization

  • This type of hybridization is characteristic of alkynes such as ethyne, C2H2
  • In its excited state, each of the two carbon atoms involved blends its 2s orbital with its 2Px orbital to form two SP hybrid orbitals which are spread apart at 180° to each other
  • Consequently, alkynes have a linear shape
  • Each of the carbon atoms uses its SP hybrid orbitals to form two sigma bonds; one with a hydrogen atom, and the second with the other carbon atom
  • The unused 2Py and 2Pz orbitals on each carbon atom then overlap with their types collaterally to form two pi bonds
  • The triple bond between two carbon atoms in an alkyne is made up of one sigma bond and two weak pi bonds that are quite easy to break
  • The presence of two weak pi bonds in an alkyne explains why alkynes are more reactive than alkenes with only one pi bond

Photo Credit: Nigerian Scholars

Effects of Lone Pairs and Bond Pairs on the Shapes of Molecules

  • The shapes of methane, ammonia, and water would have been similar but for the presence of the lone pairs of electrons on the central atoms, nitrogen and oxygen, in ammonia and water respectively
  • All four bond pairs in methane repel one another to an equal extent and the equilibrium point reached confers a tetrahedral arrangement on the methane molecule
  • In ammonia, nitrogen is surrounded by three bond pairs and one lone pair of electrons. The lone pair of electrons repels the bond pairs much more than the bond pairs repel one another.
    • This closes up the bond angle to 107° and confers a trigonal pyramidal shape on the ammonia molecule
  • In water, oxygen is surrounded by two bond pairs and two lone pairs of electrons
    • The two lone pairs repel each other vigorously and this force of repulsion is much larger than either the force of repulsion between lone pair – bond pair or that between bond pair – bond pair
    • Consequently, the net force causes the bond angle to be reduced to only 105°
    • This in turn causes the water molecule to have a V-shape or an angular shape

Shapes of Diatomic Molecules and CO2

  • Every diatomic molecule has a linear shape and a bond angle of 180°
  • Thus, each of HCl, O2, Cl2 has a linear shape
  • In carbon (IV) oxide, the carbon atom is joined by a double bond to each of the oxygen atoms
  • Mutual repulsion of the electron clouds of the double covalent bonds force them to the opposite sides of the central carbon atom
  • This results in a linear arrangement with a bond angle of 180°
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