This document discusses electronic structure and bonding. It describes how electrons fill different shells and orbitals based on principal quantum number. Maximum electrons per shell is 2n2. Ionic bonding occurs via electrostatic attraction between cations and anions to form an ionic lattice. Covalent bonding involves orbital overlap between atoms to form shared pairs of electrons and molecules. Bond strength and properties depend on whether bonding is mainly ionic or covalent.

1. Electrons and Bonding
MODULE 2
CHAPTER 5

2. Electronic Structure
Shells
 They are regarded as energy levels
 Energy increases as shell number increases
 Shell number is the principal quantum number (n)
Electrons in shells
 Electrons have properties of waves and particles
 Maximum number of electrons in shell = 2n 2

3. Electronic Structure
Maximum number of electrons
Energy level Number of electrons
1 2
2 8
3 18
4 32

4. Electronic Structure
Orbitals
 A region around a nucleus that can hold up to two electrons with opposite spins
or a high probability of finding an electron
 Each orbital (s , p , f , d) is a different shape
S orbital
 Spherical
 Each shell contains one
 The greater the shell number (n) the greater its radius

5. Electronic Structure
P orbital
 Dumbbell shape
 Each shel from n=2 contain a three (at right angles from eachother)
 Greater the shell number (n) the further it is from the nucleus
D and F orbitals
 Double dumbbell shape
 Each shell from n=3 has five d orbitals
 Each shell from n=4 has seven f orbitals

6. Electronic structure
Sub shells
 Each new shell gains a new type of orbital
 Number of orbitals increases with each orbital
 Number of electrons in each sub shell also increases
Orbitals fill in order of increasing size
 Each new type of sub shell has a higher energy
 However 4S fills before 3D

7. Electronic Structure
Electrons pair with opposite spins
 Electrons have a negative charge so repel one another, and hence have a
property called spin (shown by either an upwards or downwards arrow) to
counteract the repulsion
Orbitals with same energy are occupied singularly first
 One electron occupies each orbital before pairing starts, this prevents repulsion
until there is no other choice

8. Electronic Structure
Eectronic Configuration
 E.g. Lithium 1S 2 , 2S 1
 Can be written in shorthand using the previous noble gas’ chemical symbol and
the remaining electron sub shells
Ions
 Positive ions are cations they’re formed when atoms lose electrons
 Negative ions are anions they’re formed when atoms gain electrons

9. Electronic Structure
Blocks on the periodic table
 S block (highest energy electrons in s sub shell) left block of two groups
 P block (highest energy electrons in p sub shell) right block of six groups
 D block (highest energy electrons in d sub shell) centre block of ten groups
Ions of S and P block elements
 Highest energy sub shells lose of gain electrons
Ions of D block elements
 4S sub levels fills and emptied before 3D (as 4S has lower enegy)

10. Ionic Bonding and Structure
Ionic bonding
 It is the electrostatic attraction between positively and negatively charged ions.
 Popular cations: metals and ammonium
 Popular anions: non-metals and polyatomic ions
Ionic compounds
 Outer shell electrons from a non metal atom are transferred to the outer shell of a non
metal atom
 Cations and anions are formed
 Ions have same configuration as noble gases (full outer shell)

11. Ionic Bonding and Structure
Structure of ionic compounds
 Each ion attracts oppositely charged ions in all directions
 This results in a giant ionic lattice structure containing lots of ions
Properties of ionic compounds
 Melting and boiling points: Almost all ionic compounds are solid at room
temperature. There are very strong electrostatic forces of attraction between ions
so bonds need a lot of energy to be overcome (giving the compounds high
melting and boiling points). The greater the ionic charge the larger the strength
of the bonds and hence higher the melting point

12. Ionic Bonding and Structure
 Solubility: Many ionic compounds dissolve in polar solvents (e.g water) as the
positive and negative charges of the molecule attract the ions. The greater the
ionic charge thr stronger the bonds between the ions so compounds become less
soluble
 Electrical conductivity: ionic compounds conduct electricity when molten or
dissolved as the ionic lattic breaks down allowing ions to be mobile and carry a
charge

13. Ionic Bonding and Structure
Teeth
 The enamel on your teeth is made of the ionic compound calcium
hydroxyapatite which can be removed in acidic conditions (resulting in tooth
decay). Saliva helps neutralise the acid but it’s not always enough, so toothpaste
contains fluoride in the form of calcium fluoride which replaces lost ions by
forming fluoropatite which is stronger and more resistant than hydroxypatite

14. Covalent Bonding
Covalent compounds and molecules
 Covalent bonding is the strong electrostatic attraction between a shared pair of
electrons and the nuclei of the bonded atoms.
 It occurs between atoms in: non metals, compounds of nom metals and
polyatomic ions
 The atoms are bonded together into a single unit: a small molecule (H2), a giant
covalent structure (SiO2) or a charged polyatomic ion (NH4
+)

15. Covalent Bonding
Orbital Overlap
 A covalent bond is the overlap of atomic O rbitals, each containing one electron,
to give a shared pair of electrons
 The shared pair of electrons are attracted to the nuclei of both bonding atoms
 Bonded atoms often have the same electronic structure as a noble gas (full outer
shell)

16. Covalent Bonding
Localised
 The attraction is localised (directional) and hence acts solely between the
involved atoms. This can result in a small unit (a molecule).
 A molecule is the smallest part of a covalent compound that can exist whilst
retaining the chemical properties of a compound
Lone pairs
 Paired electrons that are not shared

17. Covalent Bonding
Multiple covalent bonds
 Two atoms share more than one pair of electrons (e.g double and triple)
 Examples include carbon dioxide (double bond) and nitrogen (triple bond)
Dative covalent bond
 A active (co-ordinate) bond is one in which the shared pair of electrons come from one atom (originally
the lone pair).
 Usually bonds to a H+ ion
 Shown in display formula as arrow from the donating atom
Average bond enthalpy
 A measurement of covalent bond strength (larger = stronger )

18. Diagrams
Ionic
Bodning
Covalent
Bonding
Display
Formula