1.5 Periodicity

• Trends in physical properties down a group
• Chemical properties of group I, II and VII elements
• Trends in properties across period three
• Atomic size, ionisation energy and electronegativity
• Applications of elements of the periodic table

• describe the trends in physical properties of elements of the periodic table
• investigate the chemical properties of elements in groups of the periodic table
• describe the trends in properties across a period
• outline applications of elements of the periodic table
• appreciate applications of various elements of the periodic table

“Why does reactivity increase down one group and decrease down another?”

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1. Trends down a group
   40 min · 2 practice tasks · quiz
   Why group I gets more reactive downwards and group VII gets less.
2. Trends across a period
   40 min · 2 practice tasks · quiz
   How properties change from sodium to argon, and why.
3. Nuclear charge, distance and shielding
   40 min · 1 practice tasks · quiz
   Every periodic trend is decided by how tightly the nucleus holds the outer electron.
4. What atomic radius means
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   Atomic radius is half the distance between nuclei of two bonded atoms of the same element.
5. Atomic radius down a group
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   Each step down a group adds a shell, so atoms get larger even though proton number also rises.
6. Atomic radius across a period
   40 min · 1 practice tasks · quiz
   Across a period atoms shrink because extra protons pull the same shells inward.
7. Ionic radius versus atomic radius
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   Positive ions are smaller than their atoms; negative ions are larger.
8. First ionisation energy
   40 min · 1 practice tasks · quiz
   First ionisation energy is the energy needed to remove one electron from a gaseous atom.
9. Ionisation energy down a group
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   First ionisation energy falls down a group because the outer electron is further out and more shielded.
10. Ionisation energy across a period
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    First ionisation energy generally rises across a period; small dips at Al and S do not hide the pattern.
11. Successive ionisation energies
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    A big jump in successive IE shows when an inner shell is broken; the jump confirms the group.
12. Electron affinity
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    Electron affinity measures how strongly an atom attracts an extra electron; halogens have the largest values.
13. Electronegativity
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    Electronegativity is the pull of an atom on shared electrons in a bond; fluorine is the highest.
14. Electronegativity trends
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    Electronegativity increases across a period and decreases down a group, for the same nuclear-pull reasons as IE.
15. Electronegativity difference and bond type
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    A large EN gap gives ionic bonding; a small gap gives covalent bonding; in between sits a polar covalent bond.
16. Melting points of group I and group II metals
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    Melting points of alkali metals fall down the group because metallic bonds weaken as atoms get larger.
17. Melting and boiling points of the halogens
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    Halogen melting and boiling points rise down the group because larger molecules have stronger intermolecular forces.
18. Melting points across period 3
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    Melting point across period 3 peaks at silicon because structure changes from giant metallic to giant covalent to simple molecular.
19. Electrical conductivity across period 3
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    Na, Mg and Al conduct; silicon is a semiconductor; P, S, Cl and Ar do not conduct.
20. Physical properties of the alkali metals
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    Group I metals are soft, low-density, silvery, stored under oil, and have characteristic flame colours.
21. Chemical reactions of the alkali metals
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    Group I metals burn in air to oxides (and peroxides for Na and K) and react with water and chlorine; reactivity increases down the group.
22. Group I in Kenya: salt and Magadi soda
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    Kenya does not mine sodium metal; it mines Na+ as coastal salt and Magadi soda ash.
23. Group II compared with group I
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    Alkaline earth metals have two outer electrons, form M2+, and are harder, higher-melting and less reactive than their group I neighbours.
24. Chemical reactions of group II
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    Magnesium burns in air and reacts with steam; calcium reacts with cold water; both form M2+ compounds.
25. Group II in Kenya: limestone, cement and fluorspar
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    Kenya meets calcium as limestone for cement, fluorspar in the Kerio Valley, and as hard-water ions.
26. Halogen displacement in depth
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    A more reactive halogen oxidises a less reactive halide ion; colour changes prove the order F > Cl > Br > I.
27. Chlorine, hydrogen halides and Kenyan water
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    Hydrogen halides become stronger acids down the group; chlorine is used to disinfect Kenyan piped water.
28. Noble gases: group 0 behaviour
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    Noble gases have full outer shells, almost no chemistry, and are used because they do nothing.
29. Period 3 oxides: basic to acidic
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    Oxides change from basic Na2O and MgO through amphoteric Al2O3 to acidic SiO2, P4O10, SO2 and chlorine oxides.
30. Period 3 hydrides
    40 min · 1 practice tasks · quiz
    Hydrides change from ionic and basic (NaH) to covalent and acidic (HCl) across period 3.
31. Predicting properties from position
    40 min · 1 practice tasks · quiz
    Given only a position on the table, you can predict size, IE, EN, metal/non-metal character and the kind of oxide.
32. Period 3 chlorides and a Kenya review
    40 min · 1 practice tasks · quiz
    Chlorides change from ionic NaCl to covalent SiCl4 that hydrolyses; Kenyan industry is a map of these periodic families.