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Chapter 4: Bonding
4.1 Ionic bonding
- An ion is a charged particle.
- The number of charges on an ion is equal to the number of electrons lost (positive) or gained (negative) by an atom.
- Metals lose electrons to form positive ions (cations); non-metals gain electrons to form negative ions (anions).
- The charge on an ion can usually be predicted from the group of the element in the Periodic Table; transition metal elements can form more than one ion.
- Common polyatomic ions include: OH–, HCO3–, NO3–, CO32–, SO42–, PO43–, NH4 +.
- Ionic compounds consist of ions held together by forces of electrostatic attraction.
- Ionic compounds are electrically neutral, as they consist of a lattice in which the total number of positive charges is balanced by the total number of negative charges. The formula of the compound is expressed as its simplest ratio, e.g. the ions Xm + and Yn– will form the compound XnYm.
- In the ionic lattice, each ion is surrounded by a fixed number of ions of the opposite charge, known as the coordination number.
4.2 Covalent bonding and 14.1 Shapes of molecules and ions
- A covalent bond is the electrostatic attraction between a pair of electrons and positively charged nuclei.
- A molecule is a group of atoms held together by covalent bonds.
- Two pairs of shared electrons = double bond.
- Three pairs of shared electrons = triple bond.
- Increasing number of bonds Þ shorter and stronger bonds.
- Lewis structures show all the valence electrons of the atoms in the molecule or polyatomic ion.
- The octet rule refers to the fact that most atoms form a stable arrangement with eight electrons in their outer shell.
- Exceptions to the octet rule include:
Less than an octet – BeCl2, BF3 (central atom very small)
Expanded octet – PCl5, SF6 (central atom from third Period or beyond).
- Dative bonds form when both the shared electrons originate from the same atom.
- Polar bonds form when the two atoms bonded together have different electronegativity values.
- VSEPR theory: the total number of charge centres determines their geometrical arrangement by maximum repulsion; the shape of the molecule then depends on the number of bonding pairs within this arrangement.
|Number of charge centres||Number of bonding pairs||Shape of molecule|
- The polarity of a molecule depends on:
- its bond polarities
- its molecular shape – whether cancellation occurs between the polar bonds.
- Hybridization occurs when different atomic orbitals mix to form new atomic orbitals for bonding.
- Sigma (σ) bonds form when atomic orbitals (s, p or hybridized) overlap along the bond axis; all single bonds are σ bonds.
- Pi (π) bonds form when p atomic orbitals overlap laterally; the electron density is concentrated above and below the bond axis.
- Double bond = one σ and one π bond.
- Triple bond = one σ and two π bonds.
- The shape of hybridized orbitals:
sp3 Þ tetrahedral
sp2 Þ planar triangular
sp Þ linear.
- The number of resonance structures that can be drawn for a molecule is the same as the number of possible positions for a double bond.
- Delocalization of π electrons ® greater stability and bonds of intermediate length and strength.
4.3 Intermolecular forces
- The forces between molecules are largely determined by the charge separation within the molecule:
- Non-polar molecules Þ van der Waals’ forces
- Polar molecules Þ dipole–dipole attraction
- Polar molecules in which H is bonded to O, N or F Þ hydrogen bonding.
- In order of strength:
- van der Waals’ forces < dipole–dipole < hydrogen bonding.
- The stronger the intermolecular force, the lower the volatility (higher boiling point).
4.4 Metallic bonding
- Metal atoms are held together by the electrostatic attraction between a lattice of positive ions and delocalized electrons.
- The properties of metals – electrical and thermal conductivity, malleability, ductility – are a result of the delocalised electrons.
4.5 Physical properties
|Covalent substances||Ionic compounds|
|Solubility in polar solvents||low||high|
|Solubility in non-polar solvents||high||low|