AS Chemistry – Atomic structure and bonding

Sub-atomic particles

• Protons – mass 1; charge +1

• Electrons – mass 1/1840; charge –1

• Neutrons – mass 1; charge 0

• Atomic number – number of protons in nucleus

• Mass number – protons + neutrons in nucleus

• Electrons – (atomic number – charge)

Relative masses

• Mass relative to 1/12th mass of 1 atom carbon-12

• Relative atomic mass: weighted mean of relative isotopic masses

• Relative molecular mass:sum of relative atomic mass in molecular formula

• Empirical – simplest ratio

• Molecular – atoms in a molecule

Mass spectroscopy

• Vaporisation

• Ionisation – E (g) + e- (high energy) E+ (g) + 2e-

• Acceleration – from +; through – grid

• Deflection – magnetic field

• Detection

How many protons, neutrons and electrons

•24He

2 protons, 2 neutrons, 2 electrons•

2145Sc

21, 24, 21•

1735Cl

17, 18, 17•

1737Cl-

17,20,18

Ionisation energy & electron affinity

• 1st ionisation energy: E (g) E+ (g) + e-

• 2nd ionisation energy: E+ (g) E2+ (g) + e-

• 1st electron affinity: E (g) + e- E- (g)

• 2nd electron affinity: E- (g) + e- E2- (g)

Electron configuration• Sequential ionisation energies: major quantum

shells

0

2 0

4 0

6 0

8 0

1 00

1 2 0

1 4 0

1 6 0

1 8 0

2 00

1 2 3 4 5 6 7 8 9 10 11 12

• Sodium 2, 8, 1

Electron configuration• 1st ionisation energies: quantum sub-shells

0

5 0

1 00

1 5 0

2 00

2 5 0

H He Li Be B C N O F Ne Na

• Sodium 1s2 2s2 2p6 3s1

Electron filling

• Animation

Write full electron configurations for:

•6C 1s22s22p2

•17Cl 1s22s22p63s23p5

•20Ca 1s22s22p63s23p64s2

•23V 1s22s22p63s23p64s23d3

•29Cu 1s22s22p63s23p64s13d10

Bonding

• Ionic/electrovalentopposite charges attracthigh m.p./b.p. – a lot of energy needed

to overcome attractionsconducts as liquid/solution because ions

can move• Covalent

Shared electron pairs – one from each atom• Dative/coordinate covalent Shared electron pairs – both from same atom

Type of covalent bond

• Horizontal overlap of orbitals:σ bond

Type of covalent bond

• Perpendicular overlap of orbitals:π bond

Limited rotation about π bond can result in cis-trans isomerism

Hydrated metal ions

Mn+

OHH

OH

H

OH

HO

HH

OHH

OH

H

Dative bond

Covalent bond

Metallic bonding

Sodium chloride – in the crystal each sodium ion forms 6 ionic bonds to adjacent chloride ions

Electron density map for hydrogen molecule

High concentration of negative charge between H nuclei.

This is strongly attracted by both nuclei so attractive interactions exceed repulsive ones

Bonding determines structure • Extensive bonding in all directions in

space results in giant structures• Metals and ionic compounds always

have giant structures • Most covalent compounds exist as

molecules but some have giant structures eg diamond and silica (SiO2)

• If no bonds are formed then the substance is a monatomic gas

Intermediate nature of bond• Electronegativity

Atoms at opposite end of bond have differentattraction for bond pair of electrons

More electronegative atom becomes slightly -ve

Hδ+-Clδ-

• Polarising

Small highly charged cations polarise anionstowards covalency

1+ 2+

Polar bonds and polar molecules

• Symmetrically arranged polar bonds produce a non-polar molecule

Cl

Cl

C

Cl

Cl

H

Cl

C

Cl

Cl

PolarNon-polar

Intermolecular forces

• Van der Waals

Moving electrons produce a temporary dipole

Temporary dipoles induce dipoles

Size of dipole proportional to number of electrons

Closer packing gives greater attraction

Intermolecular forces

• Permanent dipole

Polar molecules have stronger attractions

• Hydrogen bonds

Highly electronegative atom (& small)

Lone pair of electrons

Bonded to hydrogen

Hydrides

• General trend

Higher mass – larger van der Waals forces• Period 2

Van der Waals forces + hydrogen bonds

-40

-20

0

20

40

60

80

100

120

2 3 4 5

Giant covalent (network covalent)

• Lattice of atoms joined by covalent bonds

• Lots of energy required to separate covalently bonded atoms

• Insulator – electrons fixed

• Graphite: covalently bonded layers (σ bond) delocalised electrons between layers (π bond)

Simple molecular

• Strong covalent bonds between atoms

• Low melting & boiling point

Weak forces (van der Waals/permanent dipole/hydrogen bonds) between molecules

• Insulator – no charged particles available to move

Giant Ionic Lattice

• Oppositely charged ions

• Lots of energy required to overcome strong electrostatic attraction

• Insulator as solid – ions fixed

• Conducts as liquid/solution – ions able to flow

Molecular shape

• Valence Shell Electron Pair Repulsion Theory

• Electron pairs in outer shell of central atom as far apart as possible in order to minimise repulsion

• Include bonded and non-bonded (lone) pairs

Shapes• 1 pair – linear

H-H• 2 pairs – linear

Cl-Be-Cl• 3 pairs – trigonal planar

Cl

BClCl

• 4 pairs – tetrahedral

Cl

Cl

C

Cl

Cl

• Lone pairs act as ‘invisible’ bonds

..

SOO

bent

..

H

O

:

H

..

H

N

H

H

..

H

F

:

..

Pyramidal Bent Linear

•Repulsion:

lone pair:lone pair > lone pair:bond pair > bond pair:bond pair

Bond Angles

107o 105o