A-Level Chemistry notes that cover everything students need to get top grades- such as key practicals and their results, as well as how to show these, models that students need to recognise in the exam, as well as tricky key words and balanced equations students need.
The Structure of the Atom
Mass Spectrometry
Electronic Structure
Ionisation Energies
, THE STRUCTURE OF THE ATOM
a) Protons, neutrons and electrons
Atoms are made up of three fundamental particles: protons, neutrons and electrons.
Protons and neutrons are found in the nucleus and are collectively called nucleons.
Electrons orbit the nucleus in a similar way to that in which planets orbit a sun. In
between the electrons and nucleus there is nothing (empty space).
The nucleus is very small; if an atom were the size of a football pitch, the nucleus
would be the size of a drawing pin.
The basic properties of these three particles can be summarized in the following table:
Particle Charge Mass
Proton +1 unit Approx 1 unit
Neutron No charge Approx 1 unit
Electron -1 unit Approx 1/1840 units (very small)
1 unit of charge is 1.602 x 10-19 coulombs. A proton is given a charge of +1 and an
electron a charge of -1. All charges are measured in these units.
1 unit of mass is 1.661 x 10-27 kg. This is also not a convenient number, so we use
“atomic mass units”.
Since the mass of protons and neutrons varies slightly depending on the nucleus, then
in order to define an “atomic mass unit” we need to choose one nucleus as a standard.
For this purpose 126C , or “carbon-12”, was chosen because its mass per nucleon
(1.661 x 10 –27 kg) is around average, which means all the other nuclei have masses
close to whole numbers. An atomic mass unit is thus defined as 1/12th of the mass
of one atom of carbon-12. Everything else is measured relative to this quantity.
b) Atomic numbers, mass numbers and isotopes
An atom is named after the number of protons in its nucleus. If the nucleus of an atom
has 1 proton, it is hydrogen; if it has two protons, it is helium; if it has 3, it is lithium
etc. The number of protons in the nucleus of an atom is called the atomic number. It
has the symbol Z.
The atomic number is the number of protons in the nucleus of an atom
,Not all atoms of the same element have equal numbers of neutrons; this may vary
slightly. The sum of the number of protons and neutrons in the nucleus of an atom is
called its mass number. It is represented by the symbol A.
The mass number is the sum of the number of protons and neutrons in the nucleus of
an atom
The nucleus of an atom can thus be completely described by its mass number and its
atomic number. It is generally represented as follows:
A
ZE
Eg. 94Be, 126C, 2412Mg
Atoms with the same atomic number but with different mass numbers (ie different
numbers of neutrons) are called isotopes.
Isotopes are atoms with the same atomic number but with different mass numbers
In a neutral atom, the number of protons and electrons are the same. However, many
elements do not exist as neutral atoms, but exist as ions. Ions are species in which the
proton and electron numbers are not the same, and hence have an overall positive or
negative charge. The number of electrons in a species can be deduced from its charge:
Eg
24 2+
12Mg : 12p, 12n, 10e
24 +
12Mg : 12p, 12n, 11e
24
12Mg 12p, 12n, 12e
24 -
12Mg : 12p, 12n, 13e
, c) Relative atomic mass
The mass of an atom is measured in atomic mass units, where one unit is 12th of the
mass of one atom of carbon-12.
The relative isotopic mass of an isotope is the ratio of the mass of one atom of that
isotope to 1/12th of the mass of one atom of carbon-12.
It is usually very close to a whole number ratio:
Isotope Mass number Relative isotopic mass
1
1H 1 1.006
4
2He 4 4.003
9
4Be 9 9.012
27
13Al 27 26.919
59
27Co 59 58.933
The masses of protons and neutrons vary slightly from isotope to isotope, so the
relative isotopic mass is not exactly a whole number.
The relative atomic mass of an atom is the ratio of the average mass of one atom of
that element to 1/12th of the mass of one atom of carbon-12.
The RAM is the average mass of all the isotopes, and is often not close to a whole
number:
Some elements and compounds exist as molecules; these also have a characteristic
mass:
The relative molecular mass of a molecule is the ratio of the average mass of that
molecule to 1/12th of the mass of an atom of carbon-12.
The relative molecular mass of a molecule is the sum of the relative atomic masses of
its constituent atoms.
Eg The relative molecular mass of CO2 is 12.0 + 16.0 + 16.0 = 44.0
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