Chapter 5: Electrons and Bonding

Shells

Maximum number of electrons in each shell = 2(n^2), where n = shell number

Energy increases with shell number

The shell number is called the principal quantum number, n

Atomic orbitals

Orbital: a region of space around the nucleus that can hold up to 2 electrons with opposite spins

Models: visualise an orbital as a region of space where there is a high probability of finding an e-

Electron cloud: a negative-charge cloud with the shape of the orbital

Orbitals fill up singlely before pairing up

S-orbitals

• Electron cloud is within the shape of a sphere
• Each shell, from shell 1, contains 1 s-orbital
• Radius increases with shell number

P-orbitals

• Electron cloud is within the shape of a dumbbell
• Each shell, from shell 2 contains 3 p-orbitals
• The greater the shell number, the further the p-orbital is from the nucleus
• The 3 separate p-orbitals at right angles to each other: p(x), p(y) and p(z)

D-orbitals and F-orbitals

• Each shell, from shell 3, contains 5 d-orbitals
• Each shell, from shell 4, contains 7 f-orbitals

Orbitals of the same type are grouped together as subshells; 1xs (2), 3xp (6), 5xd (10), 7xf (14)

Electron configuration of Xe: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6

Electron configuration

The 4s subshell fills before the 3d subshell because 4s has a lower energy level

• Shells
• Subshells          1s2 2s2 2p6
• Electrons

Electron configuration of ions: 4s electrons are first in and first out

Electron configuration should be written in shell order; with 3d before 4s:

1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 (Kr)

...because when an ion is formed, electrons are lost from the 4s subshell before the 3d subshell

• E.g. Nickel (Z = 28)
• Ni    = 1s2 2s2 2p6 3s2 3p6 3d8 4s2
• Ni2+ = 1s2 2s2 2p6 3s2 3p6 3d8

Ionic bonding: the strong electrostatic attraction between oppositely charged ions

The simplest ionic compounds contain a metal and a non-metal

The ions that are formed usually have the same electron configuration as the nearest noble gas

Example: NaCl is an ionic substance because there is a strong electrostatic attraction between the positively charged sodium ion and the negatively charged chloride ion

Structure: giant ionic lattice

• Each ion attracts oppositely charged ions in all directions
• Contains billions of ions
• Actual number of ions can only be determined by the size of the crystal
• Regular arrangement - each ion is surrounded by oppositely charged ions

Melting and boiling point

• Most ionic compounds have high melting and boiling points
• High temperature needed to provide the large quantity of energy required to break the strong electrostatic attraction
• The higher the charge on the ions, the higher the Mpt / Bpt because stronger attraction
• The larger the ions are in size (more electrons) the higher the M/Bpt because stronger attr.
• Most ionic compounds are solid at RT because there is insufficient energy to overcome the strong electrostatic attraction between ions 

Solubility

• Ionic compounds only dissolve in polar solvents e.g. H2O
• Polar molecules break down the lattice and surround each ion in solution e.g. in NaCl
  • Na+ and O2- (from water) attract
  • Cl- and H+ (from water) attract

General trend: solubililty decreases as ionic charge increases because attraction gets stronger e.g. CaCO3 (Ca2+ and CO2-) has a higher Mpt / Bpt than NaCl (Na+ and Cl-)

Electrical conductivity of ionic compounds

Solid: non-conductor

• Ions are in fixed positions within the giant ionic lattice
• So there are no mobile ions (charged particles) to carry a charge
• Can't conduct electricity; non-conductor in solid state

Liquid (molten) or aqueous: conductor

• Solid ionic lattice is broken down
• Ions are mobile and free to move through the giant ionic lattice
• Ions can carry a charge; can conduct electricity in liquid and aqueous states

Summary: ionic compounds...

• Have high melting and boiling points
• Dissolve in polar solvents
• Conduct electricity in liquid and aqueous states, not when solid

Covalent bonding: the electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

Dative covalent bond: a covalent bond in which both electrons in a shared pair are supplied by only one of the bonded atoms. Shared pair was originally a lone pair

Difference between covalent and dative covalent bond: in a covalent bond, 1 electron comes from each atom in the shared pair, in a dative covalent bond, both electrons in the shared pair come from the same atom.

Orbital overlap: each orbital contains 1 electron, the orbitals overlap creating a shared pair of electrons which is attracted to the nuclei of both atoms

A covalent bond is localised: the attraction acts solely between the shared pair of electrons and the nuclei of the bonded atoms

A molecule is the smallest part of a covalent compound that can exist whilst retaining the chemical properties of the compound

Covalent compounds: small molecules, giant covalent structures and polyatomic ions (NH4+)

Displayed formula:

•                                                 **
• H2 ---> H - H            NH3 ---> H - N - H
•                                                  I
•                                                 H

Lone pairs: pairs of electrons that are not shared / involved in the bond

The number of bonds formed depends on the number of electrons in the outer shell of the central atom

Exceptions: boron

• In BF3, boron only contains 6 electrons (3 pairs) in its outer shell
  • This means bonding can't be based solely on the noble gas electron structure
  • The 3rd electron shell can hold 18 electrons so more e- are available for bonding
  • For fluorides of phosphorus, sulfur and chloride the structure contains the same number of bonds as there are F atoms e.g. PF5 (5 bonds), SF4 (4 bonds), ClF7 (7 bonds)

Exceptions: sulfur - expanded octet

• In SF6, sulfur's 6 outer shell electrons are paired so S contains 12 electrons
• A 3d subshell becomes available for expansion

Double covalent bonds: electrostatic attraction between 2 shared pairs of electrons and the nuclei of the bonded atoms e.g. CO2 ---> O = C = O or O2 ---> O = O

Triple covalent bonds: electrostatic attraction between 3 shared pairs of electrons and the nuclei of the bonded atoms e.g. N2 ---> N = N or HCN ---> H - C = N

Dative covalent bonds in ions:

• NH3 bonds with H+ ion to form NH4+
  • N shares its lone pair of electrons with H+ forming a dative covalent bond
• H2O bonds with H+ ion to form H3O+
  • O shares one of its lone pairs of electrons with H+ forming a dative covalent bond

Average bond enthalpy: a measure of covalent bond strength: the larger the value, the stronger the covalent bond