Electrons and bonding

Electron structure

Electrons and shells

• Shells regarded as energy levels which increases as shell number increases

• Shell number is called the principle quantum number n

• Shells are made up of atomic orbitals

• An atomic orbital is a region around the nucleus which can hold up to two electrons with opposite spin (high probability of finding an electron in that region)

Types of orbitals

• Each shell from n = 1 contains an s orbital

• For an s orbital the electron cloud is in the shape of a sphere

• The greater the shell number the greater the radius of the s orbital

• In p orbitals electron cloud shape of coloured in 8

• 3 different configurations of p orbital (see textbook)

• Each shell from n = 2 contains 3 p orbitals

• The greater shell number the further the p orbital is from the nucleus

• Each shell from n = 3 contains 5 d orbitals

• Each shell from n = 4 contains 7 f orbitals

Filling of orbitals

• As the 3d subshell is higher energy than the 4s subshell, the 4s subshell fills before the 3d subshell

• Electrons are –ve charged, and the opposite spins of electrons in orbitals helps to counteract the repulsion due to their charges (spin shown with opposite arrows)

• Within a subshell, orbitals have same energy

• One electron occupies each orbital before pairing starts – this prevents any repulsion between paired electrons until no further orbital available at the same energy level

• Electron configuration of krypton: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

• Short hand notation uses the previous noble gas in square brackets eg. [He], then the remaining subshells

• For example potassium: [Ar] 4s1

Electron configuration of ions

• +ve ions (cations) formed when atoms lose electrons

• -ve ions (anions) formed when atoms gain electrons

• Periodic table can be divided into blocks corresponding to their highest energy sub-shell

• When losing electrons, 3d orbital loses its electrons before the 4s orbital

Ionic bonding

• Ionic bonding is the electrostatic attraction between positive and negative ions

• Holds together cations and anions in ionic compounds

• Ionic structures form giant ionic lattices containing billions of atoms

• Many ionic compounds dissolve in polar solvents

• The polar solvent molecules break down the lattice and surround each ion in solution

• The ionic lattice must be broken down, then the water molecules must attract and surround the ions

• Solubility of an ionic compound therefore depends on the relative strengths of the attractions within the giant ionic lattice and the attractions between the ions and water molecules

Covalent bonding

• Covalent bonding is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

• Covalent bond attraction is localised – acting solely between the shared pair of electrons and the nuclei of the 2 bonded atoms

• A covalent bond is the overlap of atomic orbitals each containing one electron to give a shared pair of electrons

• Boron has 3 outer electrons, and can form 3 covalent bonds which then gives it 6 electrons, which is considered a full shell

• A dative covalent bond or coordinate bond is a covalent bond in which the shared pair of electrons has been supplied by one of the bonding atoms only

• The shared pair of electrons was originally a lone pair of electrons

• An example is ammonium (NH4+)

• Average bond enthalpy serves as a measure of the covalent bond strength – the larger the value the stronger the bond