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):
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Relative Charge:
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Proton
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1
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+1
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Neutron
|
1
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0
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Electron
|
1/2000
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-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:
- Look up
the atomic number in the
periodic table – the atomic number is always the smallest number and tends
to be below the symbol.
- The atomic number will tell you the
number of protons and therefore
the number of electrons.
- Arrange
the electrons in ‘shells’.
e.g. The electronic structure of oxygen:
- The atomic number for oxygen is 8.
- Therefore there are 8 protons and 8 electrons.
- 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:
|
|||
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1+
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Ammonium
|
NH4+
|
1-
|
Bromide
|
Br-
|
|||
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|
Copper (I)
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Cu+
|
|
Chloride
|
Cl-
|
|||
|
|
Hydrogen
|
H+
|
|
Hydroxide
|
OH-
|
|||
|
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Potassium
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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.
