1. 2 Atomic Structure

2. Properties of subatomic particles
Name
Relative charge
Relative mass(amu)
Location
Proton +1 1
nucleus
Neutron
Electron -1
Outside nucleus

3. Terminology for the Atom
Atomic no (Z): no of protons
Mass No (A): no of protons + no of neutrons
Isoptopes: atoms of the same number of protons (the same element) but different numbers of neutrons
Atomic mass unit: 1/12 the mass of a carbon-12 atom. The mass of a carbon-12 atom is defined as exactly 12 atomic mass units
Atomic mass: the average of the masses of an elements naturally occurring isotopes weighted to their abundances

4. Isotope Calculations
Boron has 2 isotopes 10B and 11B. They are present in naturally occurring boron respectively at 18.7% and 81.3%. Calculate the relative atomic mass of boron.
Ar = (18.7 x 10) + (81.3 x 11)
100
= 10.8

5. The element copper has relative atomic mass 63
The element copper has relative atomic mass and contains atoms with mass numbers 63 and 65. What is the percentage composition of a normal isotope of copper?
65x + ((100-x) x 63) = 63.55
100
65x – 63x = 6355
2x =
x = 27.5%
100 – x = 72.5%
% composition = 27.5% 65Cu 72.5% 63Cu

7. Bonding Terminology
Ionic compounds: form when an atom of one element transfers electrons to an atom of another element
Covalent compounds: form when two atoms share electrons
Ion: a charged particle
Cation: a positively charged particle
Anion: a negatively charged particle
Monoatomic ion: an ion composed of a single aton
Polyatomic ion: two or more atoms bonded covalently and having net positive or negative charge e.g. NH4+, SO42-

8. Electronic Configuration
Electrons are present in shells around the nucleus
The first shell can hold 2 electrons, the second 8 and the third 18
The no of outer shell electrons is the same as the group no

9. Find the electronic configuration of sodium
Na atomic no = 11  there are 11 protons and 11 electrons
Electronic Configuration is 2,8,1
Find the electronic configuration of chlorine
Cl atomic no = 17  there are 17 protons and 17 electrons
Electronic configuration is 2,8,7

10. Compounds
Ionic compounds are formed between a metal and a non metal e.g. magnesium chloride
Covalent compounds are formed between two or more non-metals e.g ammonia (NH3)

11. Formation of Covalent Bonds Drawing dot and cross diagrams
Only outer shell electrons are shown
Dots and crosses used to distinguish electrons from different atoms

12. Formation of HCl o o o x o o o o H Cl o o o x o o o o
HCl

13. Draw dot and cross diagrams for methane (CH4), ammonia (NH3) and nitrogen N2 and carbon dioxide (CO2)

14. Formation of ionic bonds
Elements in Group 1 form unipositive cations e.g. Na+
Elements in Group 2 form dipositive cations e.g. Mg2+
Elements in Group 3 form tripositive cations e.g. Al 3+
Elements in Group 7 form uninegative anions e.g. Cl-1
Elements in Group 6 form dinegative anions e.g. O2-

15. o o x o
Na Cl o o o o
Na Cl-
NaCl

16. x x o o
Mg F o o o o F
Mg F-
MgF2

17. Draw diagrams to represent the ionic bonding for aluminium iodide and sodium oxide

18. Properties of Ionic Compounds
High mp/bp
Conduct electricity when molten or in aqueous solution
Dissolve in polar solvents (eg water)
Hard and brittle
React readily with each other in solution

19. Covalent Compounds & Structures
Covalent compounds may be classed as simple e.g water, ammonia, chlorine, sulphur dioxide, carbon dioxide
or as giant e.g. silicon dioxide (sand) diamond, graphite

20. Simple covalent compounds are small molecules held together by Van der Waals forces only
Giant covalent structures are giant lattices where every atom is covalently bonded to many atoms

21. Diamond Structure

22. Properties of Simple Covalent Compounds
Low mp/bp
Non conducting
Soluble in non-polar solvents
Solids are soft

23. Properties of Giant Covalent Structures
High mp/bp
Non-conducting (except graphite and some semiconductors e.g. silicon dioxide)
Non-soluble
Hard (except graphite)