Periodicity
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Some exam 'Dos and Don'ts": Periodicity
Advice on answering questions on topics often highlighted by examiners as being poorly answered in the past. Make sure that the detail covered by the question is included in your syllabus specification.
• Trends across the Periodic Table
Statement of the the trends in properties of elements across a period or down a group or comparing two elements, and their explanation, are favourite exam questions which are generally not well answered. Answers often merely state the trend, without explaining the reasons for it. If asked to compare two atoms or ions, both should be mentioned in the answer.
Almost all the properties that are asked about in exam questions rely on the attraction between the outer electrons of an atom and the nucleus. Your answer to questions on trends or differences should always include a statement about how this attraction is affected.
First remember that:
Ionisation energy refers to the energy needed to remove an electron from a gaseous atom or ion, i.e. an isolated one, not part of a solid, liquid or a molecule. It is always endothermic.
Electron affinity is NOT the opposite of ionisation, but involves the addition of an electron to an atom or ion, which can be exo- or endothermic.
Electronegativity is a property (there are several scales) which measures the attraction of an atom for the pair of outer shell electrons in a covalent bond with another atom.
Trends usually depend on three main factors:
1. The charge on the nucleus (no. of protons)
2. The energy level occupied by the outer electrons and their number.
3. The shielding of outer electron from the attraction of the nucleus by underlying full shells, and repulsion between them.
One, or more, of these will be the determining factor in whether the overall attraction between the outer electrons and the nucleus increases or decreases as you move around the Periodic Table.
• Trends down a Group
Exams questions will be worth 1-4 marks, depending on whether an explanation of the trend is required. Here's a possible model (bracketed parts not essential if space or marks are limited):
•(Although nuclear charge increases), the (same number of) outer electron(s) occupy a new shell/main energy level) further from the nucleus each time.
•Completed inner shells shield the outer electrons from the nuclear attraction.
•Thus the outer shell electrons are further and further from nucleus and are less strongly held.
•Atomic radius increases, ionisation energy and electronegativity decrease.
For metals, which tend to lose electrons in reactions, reactivity increases (as ionisation energy decreases). For non-metals, which tend to gain electrons, the trend in reactivity is the opposite - compare Groups 1 & 7.
Metal ions tend to be much smaller than the corresponding atom as the complete outer shell is often lost, and the fewer remaining electrons experience a stronger attraction by the nucleus.
Trends across a Period
•Both the nuclear charge and the number of electrons steadily increases
but the electrons all occupy the same main energy level (outer shell).
•(As no new shells are added, the shielding by the completed shells is little changed).
•The outer electrons are thus more strongly held by the nucleus and are pulled closer.
•Atomic radius decreases but ionisation energy and electronegativity increase.
Note: although there is a general increase in ionisation energy across a period, closer examination of the values will show small drops in I.E. for Al and S compared to Mg and P respectively. Some exam questions may ask you to explain this, and answers are usually poor.
Al is explained in terms of the type of electron being lost, from a higher energy (therefore less energy to remove it) 3p-orbital instead of a 3s-orbital for Mg. S has one p-orbital filled with two electrons, whereas P has singly filled p-orbitals. Repulsion between the electrons in the filled orbital makes one of them easier to remove.
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updated: 12 May 2007
