IB DP Chemistry Topic 3.1 Periodic table SL Paper 2

a. \({A_{\text{r}}} = \frac{{[(0.56 \times 84) + (9.90 \times 86) + (7.00 \times 87) + (82.54 \times 88)]}}{{100}}\);

\( = 87.71\);

Award [1 max] if answer not given to two decimal places.

Award [2] for correct final answer.

Apply –1(U) if answer quoted in g or g\(\,\)mol^–1.

first ionization energy: \({\text{M(g)}} \to {{\text{M}}^ + }{\text{(g)}} + {{\text{e}}^ – }{\text{/e}}\) / the (minimum) energy (\({\text{in kJ}}\,{\text{mo}}{{\text{l}}^{ – 1}}\)) to remove one electron from a gaseous atom / the energy required to remove one mole of electrons from one mole of gaseous atoms;

b.ii. periodicity: repeating pattern of (physical and chemical) properties;

2.8.8/sp version;

b.iv. Accept any two of the following:

the outer energy level/shell is full;

the increased charge on the nucleus;

great(est) attraction for electrons;

17 p in Cl nucleus attract the outer level more than 11 p in Na nucleus / greater nuclear charge attracts outer level more;

Allow converse for Na.

Do not accept larger nucleus.

\({{\text{S}}^{2 – }}\) has one proton less/ smaller nuclear charge so outer level held less strongly / OWTTE;

Allow converse for chloride.

Do not accept larger nucleus.

the radii of the metal atoms increase (from \({\text{Li}} \to {\text{Cs}}\)) (so the forces of attraction are less between them) / OWTTE;

b.viii.

the forces of attraction between halogen molecules are van der Waals;

forces increase with increasing mass/number of electrons;

Detailed Solution:

a. Relative atomic mass = Sum of products of isotope mass and fraction abundance

Relative atomic mass of Sr = 84*.0056+86*.0990+87*.07+88*.8254 = 87.71

b.

i.

The first ionization energy is defined as the energy required to remove the outermost (most loosely held) electron from a neutral, gaseous atom to form a positively charged ion (cation). It is a measure of the strength of attraction between the nucleus and the valence electron(s) of an atom.

Mathematically, the first ionization energy can be represented by the following equation:

X(g) → X^+(g) + e^-

ii.

Periodicity refers to the repeating pattern of properties and behaviors of elements in the periodic table. It is the fundamental concept that underlies the organization of the elements based on their atomic number and electron configuration.

The periodicity of elements arises from the periodic variation of their electron structure, which determines their chemical and physical properties. Elements within the same group (column) of the periodic table share similar characteristics due to having the same number of valence electrons, which determines their reactivity and chemical behavior.

iii.

In order to write the Argon electron configuration we first need to know the number of electrons for the Ar atom (there are 18 electrons). When we write the configuration we’ll put all 18 electrons in orbitals around the nucleus of the Argon atom.

The Argon electron configuration will be 1s²2s²2p⁶3s²3p⁶.

iv. The noble gases helium, neon, and argon have the highest first ionization energies in their respective periods due to their stable and fully filled electron configurations, making it energetically unfavorable to remove an electron from them.

v.

Chlorine (Cl) has a smaller atomic radius than sodium (Na) due to two main factors: effective nuclear charge and the number of electron shells.

Effective nuclear charge refers to the attractive force experienced by the outermost electrons of an atom from the nucleus. In chlorine, the effective nuclear charge is higher than in sodium. Although both elements have the same number of electron shells (3), the atomic number of chlorine (17) is higher than that of sodium (11), meaning it has more protons in the nucleus. The increased positive charge in the nucleus of chlorine attracts the electrons more strongly, resulting in a smaller atomic radius compared to sodium.

Additionally, the number of electron shells is the same for both chlorine and sodium, but chlorine has more electrons in its outermost shell. Chlorine has seven valence electrons in its 3rd energy level (3s^2 3p^5), while sodium has only one valence electron in its 3rd energy level (3s^1). The increased electron-electron repulsion in the outermost shell of chlorine causes the electrons to be slightly pushed away from each other, leading to a smaller atomic radius compared to sodium.

vi.

The sulfide ion (S2-) is larger than the chloride ion (Cl-) due to the difference in the effective nuclear charge.

The effective nuclear charge is the net positive charge experienced by the outermost electrons of an ion. In the case of the sulfide ion, it has 16 protons in its nucleus, while the chloride ion has 17 protons. Both ions have gained two electrons to achieve a stable electron configuration. However, the sulfide ion has one less proton compared to the chloride ion, resulting in a lower effective nuclear charge. This lower effective nuclear charge leads to weaker attraction between the nucleus and the outermost electrons in the sulfide ion, causing it to have a larger atomic radius compared to the chloride ion.

vii.

The melting points of Group 1 metals decrease down the group due to atomic Size. As you move down Group 1, the atomic size or radius of the elements increases. This is because each successive element has an additional energy level of electrons. The increase in atomic size leads to weaker metallic bonding between the atoms, making it easier to overcome the attractive forces and melt the metal.

viii. The increasing melting points of Group 7 elements (halogens) down the group are primarily due to the increasing strength of van der Waals forces, specifically London dispersion forces, which are a type of intermolecular force. With larger atoms and more electrons, there is an increase in the electron cloud’s size and polarizability. The larger and more polarizable electron cloud leads to stronger London dispersion forces between the atoms or molecules. These stronger intermolecular forces require more energy to overcome, resulting in higher melting points down the group.