1. AS Chemistry – Atomic structure and bonding

2. 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)

3. Relative masses Mass relative to 1 / 12 th 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

4. Mass spectroscopy Vaporisation Ionisation – E (g) + e - (high energy)  E + (g) + 2e - Acceleration – from +; through – grid Deflection – magnetic field Detection

5. How many protons, neutrons and electrons 2 4 He 2 protons, 2 neutrons, 2 electrons 21 45 Sc 21, 24, 21 17 35 Cl 17, 18, 17 17 37 Cl - 17,20,18

6. Ionisation energy & electron affinity 1 st ionisation energy: E (g)  E + (g) + e - 2 nd ionisation energy: E + (g)  E 2+ (g) + e - 1 st electron affinity: E (g) + e -  E - (g) 2 nd electron affinity: E - (g) + e -  E 2- (g)

7. Electron configuration Sequential ionisation energies: major quantum shells Sodium 2, 8, 1

8. Electron configuration 1 st ionisation energies: quantum sub-shells Sodium 1s 2 2s 2 2p 6 3s 1

9. Electron filling Animation

10. Write full electron configurations for: 6 C 1s 2 2s 2 2p 2 17 Cl 1s 2 2s 2 2p 6 3s 2 3p 5 20 Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 23 V 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 29 Cu 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10

11. Bonding Ionic/electrovalent opposite charges attract high m.p./b.p. – a lot of energy needed to overcome attractions conducts 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

12. Type of covalent bond Horizontal overlap of orbitals: σ bond

13. Type of covalent bond Perpendicular overlap of orbitals: π bond Limited rotation about π bond can result in cis-trans isomerism

14. Hydrated metal ions M n+ O HH O HH O HH O HH O HH O HH Dative bond Covalent bond

15. Metallic bonding

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

17. 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

18. 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 (SiO 2 ) If no bonds are formed then the substance is a monatomic gas

19. Intermediate nature of bond Electronegativity Atoms at opposite end of bond have different attraction for bond pair of electrons More electronegative atom becomes slightly -ve H δ+ -Cl δ- Polarising Small highly charged cations polarise anions towards covalency 1+ 2+

20. Polar bonds and polar molecules Symmetrically arranged polar bonds produce a non-polar molecule Cl C H C Polar Non-polar

21. 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

22. Intermolecular forces Permanent dipole Polar molecules have stronger attractions Hydrogen bonds Highly electronegative atom (& small) Lone pair of electrons Bonded to hydrogen

23. Hydrides General trend Higher mass – larger van der Waals forces Period 2 Van der Waals forces + hydrogen bonds

24. 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)

25. 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

26. 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

27. 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

28. Shapes 1 pair – linear H-H 2 pairs – linear Cl-Be-Cl 3 pairs – trigonal planar Cl B

29. 4 pairs – tetrahedral Cl C Lone pairs act as ‘invisible’ bonds.. S OO bent

30. .. H O : H H N H H H F : Pyramidal BentLinear Repulsion: lone pair:lone pair > lone pair:bond pair > bond pair:bond pair Bond Angles 107 o 105 o