GRADE 10 SCIENCE
THE ATOM
A. MODELS OF THE ATOM
5th Century BC - Leucippus of Miletus (in Turkey), along with his student Democritus spoke of
atoms for the first time (from the Greek word, atomos meaning indivisible).
Early 1800’s - John Dalton, proposed the ‘Billiard ball model’. He measured the masses of
combining substances. The essence of his theory is given by:
1. Elements are made of extremely small particles called atoms.
2. Atoms of a given element are identical in size, mass, and other properties.
3. Atoms cannot be subdivided, created, or destroyed.
4. Atoms of different elements combine in simple whole-number ratios to
form chemical compounds.
5. In chemical reactions, atoms are combined, separated, or rearranged.
Dalton’s work depicted atoms as
individual spheres with a constant
density (ie billiard balls). This
model showed atoms as solid and
indivisible.
1872 - Dimitri Mendeleev devised the first periodic table - with gaps for undiscovered
elements.
1896 – Henri Becquerel, followed by Marie and Pierre Curie discover and investigate
radioactivity
1897 – JJ Thomson discovered the electron using cathode ray tube experiments. The
cathode ray tube contains a gas very low pressure, which then had a very high voltage
placed across it. The tube would glow with a stream of energy produced at the
negative electrode, the cathode.
Thomson noted that the cathode rays were deflected by both electric and magnetic
fields, deducing that the rays were a stream of negative particles of energy, called
electrons. He then passed electric current through different gases and using
different electrodes and found that the relationship of the charge to mass of the
particles making up the rays was constant.
He also likened the atom to a Plumpudding (similar to a fruit cake),
which was a sphere that contained
an equal number of positive and
negative charges spread evenly
throughout the sphere.
1911 – Ernest Rutherford – Gold foil experiment. Alpha (α-) particles (positive Helium nuclei)
were shot at a very thin piece of gold foil. Most α- particles went straight through the
foil, a few were deflected and even fewer bounced straight back.
Conclusions: - the atom consists mainly of open space in which the electrons spin
- at the centre of the atom is the nucleus that contains almost all the mass
of the atom.
1913 - Niels Bohr produced the ‘Bohr Model’ of the atom - electrons are arranged in shells, or
energy levels around the atom. He described:
- electrons in atoms having their ‘ground state’
and occupying certain, fixed orbits or quantum
levels.
- an electron could be given energy and it would
then become excited and move to a new orbit.
- on losing that energy, light would be emitted,
hence the different spectra for different
gases.
His model put scientists on the correct path, but it could only be used to explain the
spectrum for hydrogen
1924 – Louis de Broglie – an electron has wave as well has particle properties
1927 – Werner Heisenberg’s Uncertainty Principle states that we cannot know both the position
and momentum of an electron simultaneously.
1927 – Erwin Schrödinger produced the Schrödinger Wave Equation – used to mathematically
calculate a region of space in which there is a high probability of finding an electron.
 2 2

 + V(r)  (r) = E (r)

 2m

This is one form of the equation - each
letter and symbol has a significance!
1932 – James Chadwick – discovered neutrons – each one having the same mass as a proton but
no charge.
B. ATOMIC MASS AND DIAMETER
The size of an atom is incredibly small.
- one atom of Ca - diameter of 197pm – 197 picometres or 197x10-12m.
- mass of a typical calcium atom is 6.6x10-23g.
Since the mass is so small, alternative ways of looking at mass are used:
- atomic mass units ( a.m.u.) where 1 a.m.u. = 1.66 x 10-27kg
- 1 hydrogen atom has a mass of 1 amu
- relative atomic mass – one atom of carbon-12 ( 12
6 C ) was chosen as the standard, and
assigned a mass of 12 (no units). The mass of all other atoms are assigned based on
their mass relative to that of 1 atom of carbon-12.
- 1 atom of hydrogen-1 has a mass
1
12
that of carbon-12 - relative atomic mass = 1
- 1 atom of magnesium-24 has a mass 2x that of carbon 12 - relative atomic mass = 24
C. STRUCTURE OF THE ATOM
The atom consists of 3 sub-atomic particles
Sub-Atomic
Particles
Proton
Neutron
Electron
Relative mass
Relative charge
1
1
0
+1
0
-1
An atom’s structure is summarised by:
A
Z
X
A = mass number = no of nucleons
Z = atomic number = no of protons = no of electrons
X = the symbol of the element.
A – Z = no of neutrons
D.ISOTOPES
Definition: Isotopes are atoms of the same element, having the same atomic number but a
different number of neutrons.
Relative atomic mass of naturally occurring elements – this is often not a whole number (see
periodic table). Each isotope in a sample contributes to the mass, so the percentage of each
present is needed to determine the relative atomic mass.
Worked Example:
On analysis of a spectrograph, it is found that a sample of chlorine consists of 75% of the
atoms having a mass of 35 and 25% of the atoms, a mass of 37. Calculate the relative
atomic mass of chlorine.
As we are working with percentages, we can assume that from every 100 atoms, 75 have a
mass of 35 and 25 a mass of 37. The average is therefore:
35 × 75 + 37 × 25
= 35.5
100
E. ENERGY QUANTISATION and ELECTRON CONFIGURATION
Electrons occupy different energy levels in an atom, each of which has a specific energy. As
the energy of electrons is thus limited to a certain specific values, it is said to be quantised.
- higher energy levels are further away from the nucleus
- electrons in the ‘ground state’ can absorb specific amounts of
energy and move to a higher energy level (‘excited’)
- they can then drop back down to lower energy levels, thus
releasing specific amounts of energy, which is given of as
electromagnetic radiation of a specific colour.
- quantisation of energy leads to a set of specific colours for
each atom, hence atomic line spectra are observed
- Valence electrons are the electrons found in the outermost energy level of an atom.
- Core electrons are the electrons found in the innermost energy level of an atom.
- The energy levels are represented by the principle quantum
number, n, where n = 1, 2, 3, 4 etc
- The maximum number of electrons in each energy level is given
by 2n2.
- Within each energy level, electrons occupy areas in space called
atomic orbitals. We use two types, s-orbitals and p-orbitals.
- Each orbital can accommodate 2 electrons.
- within an energy level, the s-orbital has slightly lower energy
than the p orbitals.
- We only consider the orbitals shown in the diagram.
The electrons are assigned to orbitals according to the Aufbau principle, following a number of
rules:
Hunds Rule: - every orbital is singly occupied with one electron before any one orbital is doubly
occupied, and all electrons in singly occupied orbitals have the same spin
Pauli’s exclusion principle: - no two electrons in the orbital can have the same spin.
Aufbau diagrams: 1,2,3,4 = energy levels
= orbitals
 = electrons of opposite spin
4s
4s
3p
3p
3p
3s
2p
3s
2p
Potential
Energy
4s
Potential
Energy
Potential
Energy
Examples:
3s
2p
2s
2s
2s
1s
1s
1s
2-
17
Cl
N
7
16
S
Another way of representing the electron distribution is using the spdf notation (where d and f
refer to the orbitals of elements with atomic number greater than 20). Using the three examples
above:
Cl: 1s2 2s2 2p6 3s2
N: 1s2 2s2 2p3
S2-: 1s2 2s2 2p6 3s2 3p6
Shapes of orbitals
The Schrödinger Wave Equation is used to calculate the outside surface of a region of space in
which there is a high probability of finding an electron – we can thus visualise the shapes of the
orbitals: