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9.8 The Chemistry of Art: 3. Periodic Table and electron distribution

Syllabus reference (October 2002 version)
3. The distribution of electrons within elements can be related to their position in the Periodic Table	Students learn to: define the Pauli exclusion principle to identify the position of electrons around an atom identify that each orbital can contain only two electrons define the term sub-shell outline the order of filling of sub-shells identify that electrons in their ground state electron configurations occupy the lowest energy shells, sub-shells and orbitals available to them and explain why they are able to jump to higher energy levels when excited explain the relationship between the elements with outermost electrons assigned to s, p, d and f blocks and the organization of the Periodic Table explain the relationship between the number of electrons in the outer shell of an element and its electronegativity describe how trends in successive ionisation energies are used to predict the number of electrons in the outermost shell and the sub-shells occupied by these electrons	Students: process information from secondary sources to analyse information about the relationships between ionisation energies and the orbitals of electrons process information from secondary sources to use Hund’s rule to predict the electron configuration of an element according to its position in the Periodic Table

Extract from Chemistry Stage 6 Syllabus (Amended October 2002), © Board of Studies, NSW.
[Edit: 27 Jun 08]

Prior Learning: Preliminary modules 8.3.2, 8.3.3

Background: The quantum model of the atom provides the basis for the structure of the Periodic Table, where the electron configurations of the elements and the distribution of electrons within the orbitals of atoms indicate a repetitive pattern that is the basis for repeating chemical properties.

define the Pauli exclusion principle to identify the position of electrons around an atom

The Pauli exclusion principle states that no two electrons in the same atom can have the same set of four quantum numbers. That is each electron must have its own set of quantum numbers or its own unique identity. Each electron has its own 4 quantum numbers referring to – level or shell(n), sublevel or sub-shell(l), orbital(m_l) and spin(m_s)

identify that each orbital can contain only two electrons

An atomic orbital can contain zero, one or two electrons. If two electrons are present they must have opposing spins

Background information

Note that an orbital is different from an orbit.

An orbit refers to the energy level or shell in which an electron can be found. An orbit with quantum number n can contain up to 2n² electrons. The first energy level can contain up to 2 electrons, the second energy level up to 8 electrons, the third energy level up to 18 electrons, and so on.

An orbital holds a maximum of 2 electrons only.

define the term sub-shell

Energy levels are split in atoms, forming sub-shells.
Sub-shells are a consequence of two types of interactions – nucleus-electron attractions and electron-electron repulsions

outline the order of filling of sub-shells

The order of filling subshells is indicated in the two diagrams, above and below where s<p<d<f

The order in which subshells are filled can be more easily memorised by following the arrows in the diagram

process information from secondary sources to analyse information about relationships between ionisation energies and the orbitals of electrons

Useful information

Ionisation energy (IE) is the amount of energy required to remove completely one mole of electrons from one mole of gaseous atoms or ions.

The first ionisation energy (IE₁) is the energy needed to remove an electron from the highest occupied sub-shell of the gaseous atom.

Where two electrons occupy the same orbital, electron repulsions raise the orbital energy. Removing one electron relieves the repulsions and leaves a stable half filled sub-shell.

As the atomic size increases down a Periodic Table group the ionisation energy decreases, as the negative valence electron is further from the positive nucleus

The first ionisation energy generally increases across the periods of the Periodic Table due to the increase in nuclear charge. Across a period the additional electrons are going into the same shell, are about the same distance from the nucleus, but the positive nuclear charge increases.

The second ionisation energy is always greater than the first as the electron is being removed from a positive ion.

Use the information above and other information discussed in class. Analyse the information so you have a clear understanding about the relationship between ionisation energies and the orbitals of electrons.
Background Reading for Ionization Energy , The Shodor Education Foundation, Inc. Durham, North Carolina, USA.

identify that electrons in their ground state electron configurations occupy the lowest energy shells, sub-shells and orbitals available to them and explain why they are able to jump to higher energy levels when excited

When an electron is in the orbital closest to the nucleus it is in its ground state or lowest energy level.
Electrons are only able to move to a higher energy level when they absorb a photon whose energy equals the difference in energy between the two stationary states.

An interesting activity

The intense yellow lines of a sodium emission spectrum are produced when electrons that have been excited to a 3p energy level drop back to the 3s ground state level.

You can see these emission spectrum lines using a CD-ROM and a yellow sodium street light. Stand about the same distance away from a yellow street light as the light is above the ground. Hold the CD-ROM in front of you, horizontally, with the label towards the ground. The CD-ROM acts as a diffraction grating and shows the emission spectrum of sodium.

You can also see the more complex emission spectrum of mercury if you do the same with a CD-ROM near a strong white street light or sporting oval light.

process information from secondary sources to use Hund’s rule to predict the electron configuration of an element according to its position in the Periodic Table

Find information to process or use the information below.
The information below has been processed into a diagrammatic way of showing how Hund’s Rule can predict the electron configuration of elements according to their position in the Periodic Table. You may find a better way but whichever way is used, symbols make it much easier to understand.

explain the relationship between the elements with outermost electrons assigned to s, p, d and f blocks and the organisation of the Periodic Table

The Periodic Table is organised into blocks that reflect the filling of the outermost subshells with electrons.
The s block relates to two groups, Groups I and II, where Group 1 has an incomplete s-subshell and Group 2 has a complete s-subshell
The p block refers to the gradual filling of the outermost electrons in six groups III to VIII (also called groups 13 to 18).
The d block refers to the gradual filling of the outermost electrons in the transition metals.
The f block refers to the gradual filling of the outermost electrons in the lanthanide and actinide metals.

explain the relationship between the number of electrons in the outer shell of an element and its electronegativity

As the number of electrons increases in the outermost shell of an element so too does the electronegativity. Therefore, Group VII (the halogens) has the highest electronegativity as this group only requires 1 electron to complete the p subshell. Group VIII with a complete p subshell does not gain electrons.
Electronegativity usually increases from left to right across a period and usually decreases down a group.

describe how trends in successive ionisation energies are used to predict the number of electrons in the outermost shell and the sub-shells occupied by these electrons

Ionisation energy (IE) is the amount of energy required to remove an electron from a gaseous atom or ion. The first electron removed describes the the first IE, the second electron removed describes the second IE etc.
Atoms with low first ionisation energies tend to form cations during reactions while those with high first ionisation energies tend to form anions.
As the number of electrons in the outermost shell increases across a period so too does the first ionisation energy, therefore a high first ionisation energy indicates a shell or sub-shell that is almost full.

Chemistry

Extract from Chemistry Stage 6 Syllabus (Amended October 2002), © Board of Studies, NSW. [Edit: 27 Jun 08]

Extract from Chemistry Stage 6 Syllabus (Amended October 2002), © Board of Studies, NSW.
[Edit: 27 Jun 08]