Arrangement and Trends Exam Questions

(a) Elements are arranged in order of increasing atomic number (number of protons) [1]

(b) A group is a vertical column — elements in the same group have the same number of outer shell electrons and similar chemical properties [1]. A period is a horizontal row — elements in the same period have the same number of electron shells [1].

Examiner note: Mendeleev arranged elements by atomic mass, but the modern table uses atomic number. The link between group number and outer electrons is essential for predicting properties.

(a) Group III [1], Period 3 [1]

(b) Metal [1] (aluminium — elements on the left of the periodic table are metals)

Examiner note: Group number = number of outer shell electrons (for main groups). Period number = number of electron shells. Elements with 1-3 outer electrons are typically metals; elements with 5-7 are typically non-metals.

(a) Atomic radius increases down Group I [1] because each element has one more electron shell / the outer electrons are further from the nucleus [1]

(b) Atomic radius decreases across Period 3 [1] because the number of protons increases (greater nuclear charge) while electrons are added to the same shell, so the outer electrons are pulled closer to the nucleus [1]

Examiner note: Down a group: more shells = larger atom. Across a period: more protons but same number of shells = stronger pull on outer electrons = smaller atom. These are the two fundamental trends.

• Elements in the same group have the same number of outer shell / valence electrons [1]
• Chemical properties depend on the number and arrangement of outer shell electrons [1]
• They react in similar ways because they lose, gain, or share the same number of electrons [1]

Examiner note: For example, all Group I elements have 1 outer electron and all lose this electron to form 1+ ions. All Group VII elements have 7 outer electrons and all gain 1 electron to form 1- ions. Same outer electrons = same chemistry.

(a) Metallic character decreases from left to right across the period / elements change from metals to non-metals [1]

(b) Silicon has a giant covalent structure (like diamond), which requires a large amount of energy to break the many strong covalent bonds [1]

(c) They are non-metals with no delocalised electrons or free ions to carry charge [1]

Examiner note: Na, Mg, Al are metals (conduct, high mp). Si is a metalloid with a giant covalent structure. P, S, Cl are non-metals with simple molecular structures (low mp, do not conduct).

• Electronegativity is the ability of an atom to attract the bonding pair of electrons in a covalent bond [1]
• Electronegativity increases across a period from left to right (excluding noble gases) because the nuclear charge increases while the atomic radius decreases [1]

Examiner note: Fluorine is the most electronegative element. Noble gases are excluded because they do not typically form covalent bonds. The trend is driven by increased nuclear charge pulling the bonding electrons more strongly.

(a) MgO [1] (Mg is in Group II — forms Mg2+ — oxide is O2-, so MgO)

(b) KCl [1] (K is in Group I — forms K+ — chloride is Cl-, so KCl)

(c) Non-metal [1] (fluorine — Group VII elements are non-metals)

Examiner note: The group number tells you the valency for metals (Group I = 1, Group II = 2, Group III = 3). For non-metals, the valency is usually 8 minus the group number. This allows you to predict formulae.

• Metals react by losing electrons — across the period, the nuclear charge increases, making it harder to lose outer electrons, so metallic reactivity decreases [1]
• Non-metals react by gaining electrons — across the period, the nuclear charge increases and atomic radius decreases, making it easier to gain electrons, so non-metallic reactivity increases [1]
• The transition from reactive metals on the left to reactive non-metals on the right reflects the change from electron loss to electron gain [1]

Examiner note: Na is the most reactive metal in Period 3; Cl is the most reactive non-metal. The switch occurs around Si (metalloid). Reactivity depends on how easily atoms lose or gain electrons to achieve a stable configuration.