IB Chemistry SL 3.1 The Periodic Table Questions

Practise electron configurations, periodic trends, group reactions, oxide chemistry and oxidation-state reasoning across the periodic table.

Exam points

Infer group, period and s/p/d/f block from electron configurations; classify elements from periodic position.
Compare periodic trends in atomic/ionic radius, ionization energy, electron affinity and electronegativity.
Explain periodic-trend directions and exceptions using nuclear charge, shielding, sublevels and electron pairing.
Explain Group 1 reactivity with water down the group using electron loss, equations and observations.
Compare Group 17 reactivity and predict halide displacements with ionic equations and colour evidence.
Classify Period 3 oxides as basic, amphoteric or acidic using pH and acid-base reactions.
Write balanced equations for Group 1/2 and carbon or sulfur oxides reacting with water.
Calculate oxidation states in compounds and ions, including average values in mixed-valence Fe3O4.
Use oxidation-state changes to identify redox, formulas or systematic names with Roman numerals.

Covered syllabus points

3.1.1—Organization 3.1.2—Electron configuration and position 3.1.3—Periodicity 3.1.4—Group trends 3.1.5—Metal/non-metal continuum 3.1.6—Oxidation states

[Maximum number: 1]

Lithium and boron are elements in period 2 of the periodic table. Lithium occurs in group 1 (the alkali metals) and boron occurs in group 3. Isotopes exist for both elements.

Distinguish between the terms group and period.

[Maximum number: 5]

Iron rusts in the presence of oxygen and water. Rusting is a redox process involving several steps that produces hydrated iron(III) oxide, $\mathrm{Fe}_{2} \mathrm{O}_{3} \bullet \mathrm{nH}_{2} \mathrm{O}$ , as the final product. The half-equations involved for the first step of rusting are given below.

Half-equation 1: $\quad \mathrm{Fe}(\mathrm{s}) \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-}$

Half-equation 2: $\quad \mathrm{O}_{2}(\mathrm{aq})+4 \mathrm{e}^{-}+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow 4 \mathrm{OH}^{-}(\mathrm{aq})$

(a)

Identify the oxidation number of each atom in the three species in half-equation 2.

Figure for Question (a) — IB Chemistry SL

[ 2 ]

0 | $\mathrm{I} \mid-\mathrm{II}$ - II -

Marking guidance:

Award [2] for five correct.
Award [1] for four correct.
Accept use of oxidation states ( 0,+1,-2,-2,+1 ) for oxidation numbers.
Penalize once for incorrect notation (eg, 2, 2-).

(b)

State the property that determines the order in which elements are arranged in the periodic table.

[ 1 ]

(c)

State the relationship between the electron arrangement of an element and its group and period in the periodic table.

[ 2 ]

[Maximum number: 5]

Chlorine undergoes many reactions.

(a)

State, giving a reason, whether the chlorine atom or the chloride ion has a larger radius.

[ 1 ]

$\mathrm{Cl}^{-}$ AND more «electron-electron» repulsion

Marking guidance:

Accept $\mathrm{Cl}^{-}$ AND has an extra electron.

(b)

Outline why the chlorine atom has a smaller atomic radius than the sulfur atom.

[ 2 ]

Cl has a greater nuclear charge/number of protons/ $\mathrm{Z}_{\text {eff }}$ «causing a stronger pull on the outer electrons»

same number of shells

OR

same «outer» energy level

OR

similar shielding

(c)

2.67 g of manganese(IV) oxide was added to $200.0 \mathrm{~cm}^{3}$ of $2.00 \mathrm{moldm} \mathrm{m}^{-3} \mathrm{HCl}$ .

\mathrm{MnO}_{2}(\mathrm{~s})+4 \mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{Cl}_{2}(\mathrm{~g})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})+\mathrm{MnCl}_{2}(\mathrm{aq})

[ 2 ]

(i)

State the oxidation state of manganese in $\mathrm{MnO}_{2}$ and $\mathrm{MnCl}_{2}$ .
$\mathrm{MnO}_{2}:$ $\mathrm{MnCl}_{2}$ :

[ 2 ]

$\mathrm{MnO}_{2}:+4$ $\mathrm{MnCl}_{2}:+2$

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Covered syllabus points

Question 1

Question 2

Question (a)

Question (b)

Question (c)

Question 3

Question (a)

Question (b)

Question (c)

Question (i)