State what is meant by the term periodic property.
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a physical property of elements that shows a repeating pattern; when elements are arranged in order of increasing atomic number
Periodicity Melting Points And Ionisation Energies
Suggest why first ionisation energies for elements 1–36 are often plotted on a logarithmic scale.
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the values span a very wide range; a log scale compresses the range so trends and small dips within each period are clearly visible
Explain how a graph of first ionisation energy against atomic number for elements 1–36 demonstrates that ionisation energy is a periodic property.
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the graph shows a repeating pattern across each period; with peaks at noble gases (He, Ne, Ar, Kr); and troughs at Group 1 metals (Li, Na, K)
State the equation, including state symbols, that represents the first ionisation energy of magnesium.
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Mg(g) → Mg⁺(g) + e⁻; correct state symbols (g) on both Mg and Mg⁺
Explain why the melting temperature of silicon is much higher than that of phosphorus.
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silicon has a giant covalent structure with many strong covalent bonds; phosphorus exists as simple molecular P₄ with weak London forces between molecules; melting silicon requires breaking covalent bonds whereas melting phosphorus only requires overcoming London forces
Explain why the first ionisation energy of sulfur is lower than that of phosphorus.
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in phosphorus each 3p orbital contains one electron; in sulfur one 3p orbital contains two paired electrons; the paired electrons repel each other, making one easier to remove
Explain why the first ionisation energy of potassium is lower than that of sodium.
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the outer electron of potassium is in a higher principal energy level / further from the nucleus; potassium has more inner shells providing greater shielding; these effects outweigh the increased nuclear charge, so less energy is needed to remove the outer electron
Discuss the trends in first ionisation energy across Period 3, using your knowledge of electron configuration. In your answer you should: explain the general increase from Na to Ar; explain the specific dips at Al and S; compare the magnitude of these effects with the trend down Group 1.
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across Period 3, nuclear charge increases while electrons enter the same shell; shielding remains approximately constant so the outer electron is held more strongly, giving a general increase; at Al, the outer electron occupies the higher-energy 3p orbital rather than 3s, so it is easier to remove than the 3s electron in Mg; at S, two paired electrons in one 3p orbital repel each other, so one is more easily removed than the unpaired 3p electron in P; down Group 1, first ionisation energy decreases because the outer electron is in a higher shell with more shielding; this down-group decrease is much larger than the dips at Al and S, because shell number changes have a greater effect than sub-shell energy or pairing repulsion within one period