## Friday, 5 April 2013

### ATOMIC STRUCTURE AND BONDING

1.8      recall that atoms consist of a central nucleus, composed of protons and neutrons,             surrounded electrons, orbiting in shells

Self explanatory.

1.9      recall the relative mass and relative charge of a proton, neutron and electron

 Relative Mass (amu): Relative Charge: Proton 1 +1 Neutron 1 0 Electron 1/2000 -1

1.10    understand the terms atomic number, mass number, isotopes and relative atomic       mass (Ar)

Atomic Number: The number of protons in the nucleus of an atom of an element.

Mass Number: The sum of the number of protons and neutrons in an atom (measured in amu).

Isotopes: Atoms with the same number of protons but a different number of neutrons.

Relative Atomic Mass (Ar): The weighted average mass of an atom of an element on the scale on which the isotope C-12 (carbon 12) is taken to weigh 12 units exactly.

Isotopic Mass: The mass of an atom of an isotope on the scale on which an atom of C-12 is taken to weigh 12 units exactly.

1.11    calculate the relative atomic mass of an element from the relative abundances of its    isotopes

Ar = (mass number x %) + (mass number x %)
100

e.g. Bromine consists of two isotopes of isotopic masses 79 amu and 81 amu which occur in equal abundance (50% each).

Ar of Bromine = (79 x 50) + (81 x 50)
100

1.13    deduce the electronic configurations of the first twenty elements from their         positions in the Periodic Table

To work out the electronic arrangement of an atom:

1. Look up the atomic number in the periodic table – the atomic number is always the smallest number and tends to be below the symbol.

1. The atomic number will tell you the number of protons and therefore the number of electrons.

1. Arrange the electrons in ‘shells’.

e.g. The electronic structure of oxygen:

1. The atomic number for oxygen is 8.

1. Therefore there are 8 protons and 8 electrons.

1. The arrangement of the electrons will be 2,6 as the first shell only holds 2 and there’s only 6 to occupy the next shell.

1.14    deduce the number of outer electrons in a main group element from its position in the Periodic Table

The number of electrons in an element’s outer shell is the same as the element’s group number in the periodic table.

e.g. Barium is group 2 so it as 2 outer electrons.*

*Outer electrons are also known as valence electrons.

Ionic compounds

1.27    describe the formation of ions by gain or loss of electrons

Ions are atoms or molecules with an electric charge due to the gain or loss of electrons.

If electrons are lost the ion has a positive charge. Metals tend to do this so they form cations (positive ions). Elements from groups 1-3 will usually form cations.

If electrons are gained the ion has a negative charge. Non-metals tend to do this so they from anions (A Negative Ion – Anion!) So most elements in groups 5-7 will form anions.

Noble gases are inert/unreactive so do not form ions.

1.29    recall the charges of common ions in this specification

Positive Ions (Cations):                       Negative Ions (Anions):

 Charge: Name of Ion: Formula: Charge: Name of Ion: Formula: 1+ Ammonium NH4+ 1- Bromide Br- Copper (I) Cu+ Chloride Cl- Hydrogen H+ Hydroxide OH- Potassium K+ Fluoride F- Silver Ag+ Iodide I- Sodium Na+ Nitrate NO3- Lithium Li+ Hydrogencarbonate HCO3- 2+ Barium Ba2+ 2- Carbonate CO32- Calcium Ca2+ Sulphate SO42- Copper (II) Cu2+ Sulphide SO32- Iron (II) Fe2+ Oxide O2- Lead (II) Pb2+ Magnesium Mg2+ 3+ Aluminium Al3+ 3- Nitride N3- Iron (III) Fe3+ Phosphate PO43-

1.30    deduce the charge of an ion from the electronic configuration of the atom from   which the ion is formed

To find the charge of an ion look at the outer shell (valence) electrons. If the atom needs to lose an electron (e.g. Na: 2.8.1) then the ion will be positive (in sodium’s case +1). If you need to gain electrons the ion is negative.

1.31    explain using dot and cross diagrams, the formation of ionic compounds by electron   transfer, limited to combinations of elements from Groups 1, 2, 3 and 5, 6, 7

e.g. Sodium Chloride (NaCl)

Calcium Chloride (CaCl2)

1.32    understand ionic bonding as a strong electrostatic attraction between oppositely charged ions

Self explanatory.

1.33    understand that ionic compounds have high melting points and high boiling points because of strong electrostatic forces between oppositely charged ions

The strong electrostatic forces between oppositely charged ions are ionic bonds (as mentioned in 1.32). These forces require a lot of energy to break, hence high melting and boiling points.

1.34    understand the relationship between ionic charge and the melting point and boiling point of an ionic compound

In general the greater the charge, the greater the electrostatic attraction, the stronger the ionic bond and the higher the melting point.

1.35    describe an ionic crystal as a giant three-dimensional lattice structure held   together by attraction between oppositely charged ions

Self explanatory.

1.36    draw a simple diagram to represent the positions of ions in a crystal of  sodium chloride

1.47    understand an electric current as a flow of electrons or ions

Self explanatory.

1.49    understand why ionic compounds conduct electricity only when molten or in solution

Ionic compounds can only conduct electricity when their ions can move. Their ions can only move when the compound is molten or a liquid.

Covalent substances

1.37    describe the formation of a covalent bond by the sharing of a pair of electrons between two atoms

A covalent bond forms when two non-metal atoms share a pair of electrons (e.g. if one atom needs to gain two, it will share with one that needs to lose two).

1.38    understand covalent bonding as a strong attraction between the bonding pair of            electrons and the nuclei of the atoms involved in the bond

Covalent bonding is a strong attraction between the bonding pair of electrons and the nuclei of the atoms involved in the bond.

1.39    explain, using dot and cross diagrams, the formation of covalent compounds by            electron sharing for the following substances:

i           carbon dioxide (CO2)

ii           chlorine
iii         hydrogen chloride
iv         water
v          methane
vi         ammonia
vii        oxygen
viii       nitrogen
ix         hydrogen (H2)

x           ethane
xi         ethene

1.40    recall that substances with molecular structures are gases or liquids or solids with      low melting points

Self explanatory.

1.41    explain why substances with simple molecular structures have low melting points in   terms of the relatively weak forces between the molecules

Substances with simple molecular structures (see 1.4) have low melting points because of the relatively weak forces between the molecules. (All simple molecules have strong bonds between atoms but much weaker bonds between molecules).

1.42    explain the high melting points of substances with giant covalent structures in terms    of the breaking of many strong covalent bonds

Giant covalent structures are very strong because of the many covalent bonds. Giant covalent structures have high melting points because lots of strong covalent bonds must be broken.

1.43    draw simple diagrams representing the positions of the atoms in diamond     and graphite

Diamond:

Graphite:

1.44    explain how the uses of diamond and graphite depend on their structures,     limited to graphite as a lubricant and diamond in cutting

Whilst both carbon, diamond and graphite are very different. Graphite forms layers which slide over each other making a good lubricant. Diamond is very hard and is used to cut things – usually other diamonds.

1.47    understand an electric current as a flow of electrons or ions

Self explanatory.

1.48    understand why covalent substances do not conduct electricity

Covalent substances do not conduct electricity because they do not have free electrons. The exception to this is graphite.

1.45    describe a metal as a giant structure of positive ions surrounded by a sea of    delocalised electrons

Self explanatory.

1.46    explain the malleability and electrical conductivity of a metal in terms of its structure      and bonding

Metals are malleable/conduct electricity because free electrons allow metal atoms to slide over each other.