CHEMISTRY – CHAPTER 1: ATOMIC STRUCTURE
UNIT 1 : MODULE 1
1.1 Discuss the process of theoretical change with respect to Daltons Atomic Theory
Daltons Theory:
Matter consists of tiny indivisible
particles that cannot be created or
destroyed
Atoms of one element cannot be
converted into atoms of another
element
Atoms of a given element are identical
in mass and other properties and are
different from atoms of any other
element
Atom of different elements combines in
simple specific whole number ratios to
give combined atoms. This is the law of
multiple proportions.
What is now known:
Atoms are not indivisible but are
composed of subatomic particles
During a chemical reaction this is true
however in nuclear reactions atoms of
one element can be converted into
atoms of another element
Atoms of a given element can have
slight differences in their mass and
properties such as isotopes
True for most compounds but there are
a few compounds in which there are
slight variations in their atomic ratios
In order for a theory to be accepted, what criteria must be considered?
1. There must be a link between evidence and the theoretical constructs
2. Data must be reliable and accurate
3. Experiments must be replicable
4. Experiments must be a consensus within the scientific community
1.2 describe the structure of the atom;
Bohr’s Atomic Theory:
Bohr assumed that electrons move in circular orbits around the nucleus. He also
adopted the idea that energies in an atom are quantized, i.e., electrons are found in
orbits which exist at specific energy states/ levels.
Energy is emitted or absorbed by the electrons only as the electron changes from
one energy state to another. This energy is the energy emitted as or absorbed by a
photon (a particle of light). In other words, energy either is required or released
when electrons move from one state/ level to another.
Ground state – refers to when an electron is in its lowest possible energy state.
Excited state – refers to when an electron is in any other energy state besides its
ground state.
Modern Atomic Structure:
It is comprised of a small, dense, positively-charged centre called a nucleus which
contains most of its mass.
The nucleus contains protons and neutrons, while electrons can be found orbiting
the nucleus.
Protons, neutrons and electrons are called sub-atomic particles
Atomic/ Proton number (Z): Number of protons in the nucleus
Mass number (A): number of protons and neutrons in the nucleus
i.e., mass no. (A): no. of protons + no. of neutrons
An atom is neutral therefore in any atom: no. of protons = no. of electrons
Isotopes: These are atoms of the same element with the same atomic/ proton
number (Z) but different mass numbers (A), i.e., the same number of protons and
electrons but different numbers of neutrons.
The isotopes of carbon are: 11C, 12C, 13C & 14C
Sub-Atomic Particles and Electric Fields:
The effect on an electric field on beams of electrons, protons and neutrons:
Charged particles move towards the plates with opposite charge and away from the
plates with the same charge. If the same voltage to deflect electrons and protons
are used, the electrons are deflected to a much greater extent. This is because
electrons have a small mass compared to protons
Relative atomic and relative isotopic mass
The mass of a single atom is too small to weigh, so we use the 12C (carbon-12)
atom as a standard. We compare the masses of other atoms to this standard, which
has a mass of exactly 12 units.
Relative atomic mass (Ar)
Relative atomic mass (Ar) is the weighted average mass of naturally occurring
atoms of an element on a scale where an atom of 12C has a mass of exactly 12
units.
Calculating the Relative Atomic Mass (Ar) of An Element Using Its Isotopes:
Ar = ∑ (mass of each isotope x abundance of each isotope)
e.g. (1) Naturally-occurring chlorine is 75.78 % 35Cl (mass = 34.969) and 24.22 %
37Cl (mass = 36.966).
Calculate the atomic mass of chlorine:
Ar = ∑ (mass of each isotope x abundance of each isotope)
= (34.969 x 0.7578) + (36.966 x 0.2422)
= 35.45
The total abundance of all the isotopes of an element is always 100 %.
e.g. (2) Copper has an atomic mass of 63.55 and consists of two stable isotopes:
63Cu, mass = 62.930 and 65Cu, mass = 64.928. What are the percent abundances
of the two isotopes in natural copper?
Ar = (Mass of isotope 1 x Abundance) + (Mass of isotope 2 x Abundance)
63.55 = (62.930x/ 100) + (64.928(100-x)/ 100)
6355 = 62.930x + 6492.8 – 64.928x
1.998x = 137.8
x = 69
63Cu = 69 % 65Cu = 31 %
Relative isotopic mass
Relative isotopic mass is the mass of a particular isotope of an element on a scale
where an atom of 12C has a mass of exactly 12 units
Radioactivity:
Isotopes of some elements have nuclei which break down (decay) spontaneously.
These are described as radioactive isotopes. As the nuclei break down, rays or
particles are given out. These are called emissions.
The table shows three types of emission:
Uses of Radioactivity
Tracers for searching for faults in pipelines and for studying the working of certain
organs in the body. used to study thyroid function.
In medicine, for radiotherapy in the treatment of cancers.
Dating the age of objects, e.g., using 14C to date objects which were once living
Smoke detectors often use 241Am
Generating power, e.g., 235U is used in many nuclear reactors as a source of
energy.
1.3Energy Quanta and Energy Levels:
Energy Quanta and Energy Levels:
The electrons in atoms have certain fixed values of energy. The smallest fixed
amount of energy required for a change is called a quantum of energy
The electrons are arranged in energy levels which are quantised.
The atom is most stable when the electrons are in the lowest energy levels possible
for each of them. The electrons are in their ground state
When an electron absorbs a quantum of radiation it can move up to a higher
energy level. The electron is then said to be in an excited state
When an excited electron falls to a lower energy level a quantum of radiation is
given out and the electron moves to a lower energy level.
Principal quantum number (n)- describes the size of the orbital or energy level of
the atom.
• Only have positive integer values.
• As n increases, the orbital becomes larger and the electron spends
more time farther from the nucleus
Atomic Emission Spectra:
When electrical or thermal energy is passed through a gaseous sample of an
element, the radiation is emitted only at certain wavelengths or frequencies. An
emission spectrum differs from a normal visible light spectrum in that:
It is made up of separate lines.
The lines converge (get closer to each other) as their frequency
Interpreting the Hydrogen Emission Spectrum:
When electrical or thermal energy is passed through a gaseous sample of an
element, the radiation is emitted only at certain wavelengths or frequencies. Each
line in the hydrogen emission spectrum is a result of electrons moving from a
higher to a lower energy level.
Among the several series of lines seen are the:
Lyman series (seen in the ultraviolet region), where previously excited electrons fall
back to the n = 1 energy level.
Balmer series (seen in the visible region), where previously excited electrons fall
back to the n = 2 energy level.
When Ionisation Occurs:
1.4 Sub-shells and atomic orbitals
The principal quantum shells apart from the first, are split into sub-shells (sublevels). These are named s, p, d.
Atomic orbitals
Each sub-shell contains one or more atomic orbitals. An orbital is a region of space
where there is a high probability of finding an electron. Each orbital can hold a
maximum of two electrons. So the number of orbitals in each sub-shell is: s = 1, p
= 3 and d = 5.
Comparisons
Rules For Filling Electrons
The Electronic Configuration Of Atoms
Electronic Configuration of Ions
When ions are formed one or more electrons are removed from or added to the
atom. In general, the electrons are lost or gained from the outer sub-shell.
This is so because, when the 4s orbital has electrons, it has a higher energy than
the 3d orbitals
1.5 Ionisation Energies
The first ionisation energy, ΔHi1, is the energy needed to remove one electron from
each atom in one mole of atoms of an element in its gaseous state to form one mole
of gaseous ions.
Factors Influencing the Ionisation Energy Of Elements:
Size of nuclear charge: In any one period, the higher the nuclear charge (greater
number of protons), the greater the attractive force between the nucleus and outer
electrons and the more energy needed to remove these electrons. So ΔHi1 tends to
increase with increase in nuclear charge.
Distance of the outer electrons from the nucleus: The further the outer
electrons are from the nucleus, the less attractive force there is between them and
the nucleus and the lower the value of ΔHi1
Shielding: Electrons in full inner shells reduce the attractive force between the
nucleus and the outer electrons. The greater the number of full inner shells, the
greater the shielding and the lower the value of ΔHi1
First ionisation Energies Across a Period:
We can use the three ideas above to explain the change in ionisation energy with
increasing proton number:
Across a period, there is a general increase in ΔHi1. The increased nuclear charge
(number of protons) outweighs the other two factors because the successive
electrons added are being placed in the same outer electron shell (remain
constant). There is not much difference in shielding since there is the same number
of inner electron shells (remains constant).
When a new period starts, there is a sharp decrease in ΔHi1. The outer electron
goes into a shell further from the nucleus (it increases) and there is more shielding.
Both these factors outweigh the effect of greater nuclear charge. This also explains
why ΔHi1 decreases down a group.
Example:
The decrease in ΔHi1 from Be → B is due to the fact that the fifth electron in B goes
into the 2p sub-shell, which is slightly further from the nucleus than the 2s subshell. The shielding also increases slightly. These two factors outweigh the effect of
the increased nuclear charge. A similar reason explains the dip in ΔHi1 from Mg →
Al.
The decrease in ΔHi1 from N → O is due to spin-pair repulsion. The eighth electron
in O goes into an orbital which already has an electron in it. The increased
repulsion of the similarly charged particles makes. it easier to remove the eighth
electron.
Successive Ionisation Energies:
We can remove electrons from an atom one by one. The energies associated with
removing these electrons are called the successive ionisation energies.
Down a Group, I.E., decreases due to an increase in atomic radius and decrease in
Zeff (since outer electrons are further away from the nucleus and effective nuclear
charge is lower, electrons are easier/ require less energy to remove).
Past Paper Questions
May/ June 2022
May / June 2021
May/ June 2019
May / June 2018
May/
June
2017
May/ June
2016
May/ June 2015
May/June 2010
May / June 2009