2 Atomic Structure

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

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

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

The element copper has relative atomic mass 63 The element copper has relative atomic mass 63.55 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 + 6300 – 63x = 6355 2x = 6355-6300 x = 27.5% 100 – x = 72.5% % composition = 27.5% 65Cu 72.5% 63Cu

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-

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

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

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)

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

Formation of HCl o o o x o o o o H Cl o o o x o o o o HCl

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

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-

o o x o Na Cl o o o o Na+ + Cl- NaCl

x x o o Mg F o o o o F Mg2+ + 2F- MgF2

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

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

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

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

Diamond Structure

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

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