Chapter 2 “Atomic Structure” & Chapter 3 “Electrons in Atoms”
Candidates should be able to:
(a) *identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses
Sub-atomic particle
Electron
Electron (e) Mr= 1/1836 charge -1
Neutron (n) Mr= 1 charge 0
Proton (p) Mr= 1 charge +1
(b) *deduce the behaviour of beams of protons, neutrons and electrons in electric fields
(c) describe the distribution of mass and charges within an atom
(d) deduce the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers (and charge)
Nucleon number (A) – the number of protons plus neutrons in the nucleus of an atom, a.k.a. the mass number.
Proton number (Z) – the number of protons in the nucleus, a.k.a. the atomic number.
Number of neutrons – A - Z
(e) (i) describe the contribution of protons and neutrons to atomic nuclei in terms of proton number and nucleon number
(ii) distinguish between isotopes on the basis of different numbers of neutrons present
(f) *describe the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals.
A. Atomic orbital
1. A region in space where there is a high probability of finding an electron
B. Energy Levels of electrons (n)
1. Indicates the distance of the energy level from the nucleus
2. Values of n are positive integers
a. n=1 is closest to the nucleus, and lowest in energy
3. The number of orbitals possible per energy level (or "shell") is equal to n2
C. Energy Sublevels
1. Indicates the shape of the orbital
2. Number of orbital shapes allowed in an energy level = n
a. Shapes in the first four sub-shells are designated:
s=1 orbital, p=3 orbitals, d=5 orbitals, f= 7 orbitals. There can be 2 electrons in each orbital. The energy of the electrons in the sub-shells increases in the order s < p < d.
The lowest energy level will fill up first.
Which fill up in this order: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6.
Principal quantum #
( see picture above )
D. Electron Spin
1. A single orbital can contain only two electrons, which must have opposite spins
2. Two possible values for spin, +1/2, -1/2
(g) *describe the shapes of s and p orbitals
s orbitals have a spherical shell shape
p orbitals have an hourglass shape with 2 lobes
(h) state the electronic configuration of atoms and ions given the proton number (and charge)
(i) Ionization energy
(i) explain and use the term ionisation energy
Ionisation energy – is the energy needed to remove one electron from each atom in one mole of atoms of the element in the gaseous state to form one mole of gaseous 1+ ions.
(ii) explain the factors influencing the ionisation energies of elements
The energy increases for each successive removal of electrons because as each electron is removed there is a greater attractive force between the protons and the electrons. So more energy is needed to overcome this attraction.
(iii) *explain the trends in ionisation energies across a Period and down a Group of the Periodic Table (see also Section 9)
In general, the first ionization energy increases as we go from left to right across a row of the periodic table.
1. the nuclear charge increases
2. the distance between the nucleus and the outer electron remains reasonably constant.
3. The shielding by inner shells remains reasonably constant.
The force of attraction between the positive nucleus and the outer negative electrons decreases because:
1. The distance between the nucleus and the outer electron increases
2. The shielding by inner shells increases
3. These 2 factors outweigh the increased nuclear charge.
(j) deduce the electronic configurations of elements from successive ionisation energy data
Lowest values are for outer electrons, dramatic increase indicates electrons are in successive inner levels.
(k) *interpret successive ionisation energy data of an element in terms of the position of that element within the Periodic Table.
Valence electrons (outer shell) indicates what group element is in.
Extra information:
Degenerate orbitals – atomic orbitals at the same energy level. Each electron enters an orbital alone because:
1. electrons in the same area of space repel each other
2. due to their like charges
3. which minimizes repulsion.
The lowest energy arrangement is the ground-state electronic configuration, which is a description of what orbitals the atom's electrons occupy.
1. Electrons always fill the lowest energy orbitals first (Aufbau)
2. Only two electrons can occupy each orbital and they must be of opposite spin (Pauli)
3. If two or more empty orbitals of equal energy are available, place one electron in each until all are half-full (Hund)