Download YEAR 9 REVISION LIST November Exam 2013

Barnard Castle School Chemistry Department
November Exam 2013
Topic 9.1 Atomic structure and the periodic table
Structure of atoms
Atoms have a small central nucleus made up of protons and neutrons, around which
there are electrons in shells. The relative masses and charges are shown in the
table:
Name of Sub-Atomic Particle
proton
neutron
electron
Relative Mass
1
1
1/1840
(negligible)
Relative Charge
+1
0
-1
In an atom, the number of electrons is equal to the number of protons in the
nucleus. Atoms therefore have no overall electrical charge. All atoms of a particular
element have the same number of protons. Atoms of different elements have
different numbers of protons. The number of protons in an atom is called its atomic
number (proton number). When electrons, protons and neutrons were discovered in
the early 20th Century, the atoms were arranged in the modern Periodic Table in
order of their atomic number.
Isotopes are atoms with the same number of protons but different numbers of
neutrons.
For example, there are two isotopes of chlorine:
 one isotope has 18 neutrons, so when added to the 17 protons gives it a
relative atomic mass of 35. This isotope is known as chlorine-35 (35Cl)
 the other isotope has 20 neutrons, so when added to the 17 protons has a
relative atomic mass of 37. This isotope is known as chlorine-37 (37Cl).
Electrons occupy particular energy levels (also called shells). Each electron in an
atom is at a particular energy level (in a particular shell). The electrons in an atom
occupy the lowest available energy levels (innermost available shells).
Elements in the same Group (column) in the Periodic Table have the same number
of electrons in their outer shells. The arrangement of these electrons can be
drawn on diagrams, or written in a shorthand form e.g. Na is 2,8,1.
The Periodic Table
Newlands and Mendeleev attempted to classify the elements by arranging them in
order of their atomic weights. The list can be arranged in a table so that elements
with similar properties are in columns known as Groups. The table is called a
Periodic Table because elements with similar properties occur at regular intervals.
The early Periodic Tables were incomplete and some elements were placed in
inappropriate Groups if the strict order of atomic weights were followed.
Mendeleev overcame some of the problems by leaving gaps for the elements that he
thought had not been discovered.
The modern Periodic Table can be seen as an arrangement of elements in terms of
their electronic structures. Elements in the same group have the same number of
electrons in their outer shell. The trends in reactivity within Groups in the Periodic
Table can be explained because the higher the energy level (the further the outer
shell from the nucleus), the more easily electrons are lost and the less easily
electrons are gained.
Group 1: The Alkali Metals (Li, Na, K, Rb, Cs)
The elements in Group 1 of the Periodic Table are known as the Alkali Metals.
They:
 are metals with low density (the first three elements in Group 1 are less
dense than water)
 react with non-metals to form ionic compounds in which the metal ion
carries a charge of +1. The compounds are white solids that dissolve in
water to form colourless solutions
 react with water releasing hydrogen
 form hydroxides that dissolve in water to give alkaline solutions.
In Group 1, the further down the Group an element is:
 the more reactive the element
 the lower its melting and boiling point
Group 7: The Halogens (F2, Cl2, Br2, I2, At2)
The elements in Group 7 of the Periodic Table (known as the Halogens):
 have coloured vapours
 consist of molecules that are made up of pairs of atoms (diatomic)
 form ionic salts with metals in which chloride, bromide or iodide ions
(halide ions) are formed. The halide ion carries a charge of -1
 form molecular compounds with other non-metallic elements
In Group 7, the further down a group an element is:
 the less reactive the element
 the higher its melting point and boiling point
A more reactive halogen can displace a less reactive halogen from an aqueous
solution of its salt.
9.2 Bonding
Ionic bonding ONLY.
Ionic bonding happens between metals and non-metals.
METALS:
 metals have less than half full outer shells of electrons. Thus the easiest way
for them to have a full outer shell is to LOSE those outer electrons. By
losing the outer electron(s) they now have more protons than electrons, so
the particle has an overall POSITIVE charge. We say the neutral atom has
become a POSITIVE ION.
 Group 1 metals lose 1 electron and become +1 ions, e.g. Li+, Na+ and K+.
 Group 2 metals lose 2 outer electrons to become +2 ions, e.g. Mg2+ and Ca2+.
NON-METALS
 Non-metals have more than half-full outer shells of electrons. Thus the
easiest way for them to have a full outer shell is to GAIN electrons. By
gaining more electrons, they now have more electrons than protons, so the
particle has an overall NEGATIVE charge. Thus the neutral atom has become
a NEGATIVE ION.
 Group 7 elements gain 1 electron to become -1 ions, e.g. F-, Cl- and Br-.
 Group 6 elements gain 2 electrons to become -2 ions, e.g. O2- and S2-.
The positive and negative ions are attracted to each other to form an ionic bond,
which keeps the ions together as an ionic compound, for example:
 Na+ and Cl- bond to form NaCl (sodium chloride)
 Mg2+ and F- bond to form MgF2 (magnesium fluoride).
 K+ and O2- bond to form K2O (potassium oxide)
 Ca2+ and O2- bond to form CaO (calcium oxide).
You need to be able to show metals and non-metals losing or gaining electrons and
to calculate the type and size of the charge that forms.
Duration: 1 hour