This question involves knowledge of structure and bonding as well as the acidic and basic nature of oxides.

Metal oxides (MgO and Na₂O), if they react with water at all, will give alkaline solutions that therefore have a pH higher than 7.

Non-metal oxides may react with water to give acidic solutions that therefore have a pH lower than 7. However, SiO₂ (silica) is a giant covalent structure analogous to diamond, and does not react with water (sand is impure SiO₂). Sulfur trioxide, SO₃, will react (very exothermically) in water to give an acidic solution of sulfuric acid (H₂SO₄).

SO₃ is therefore the correct answer.

Group 1 metals have metallic bonding and therefore giant structures. Melting points are relatively low for metals as the atoms are large and form 1+ ions so the attraction between the nucleus and ‘sea of electrons’ is not as strong as in other metals. Nonetheless strong metallic bonds must be broken. The melting points (values in data book) become lower down the group because of the increasing atomic/ion size leads to less attraction..

Group 17 non-metals; the halogens, have covalent bonding and simple molecular structures. The atoms get bigger down the group, so the diatomic halogen molecules get bigger down the group. These non-polar molecules have weak London dispersion forces that must be broken. The melting points (data book) become higher down the group as London dispersion forces get stronger as the number of electrons/molar mass of a molecule increases.

Thus, Decrease down group 1 and increase down group 17 is the correct answer.

Metals tend to form basic (alkaline) oxides, not acidic oxides, so statement 1 is false.

Metals' outer electrons are delocalised and they have giant structures. Metals are often: good conductors of electricity and heat; high in melting and boiling point; malleable; ductile; sonorous; lustrous (shiny). Therefore statement 2 is true.

Metals have low electronegativity and tend to lose electrons when they react. The trend in ionisation energy across a period (from metals to non-metals) is a general increase. Therefore statement 3 is true.

Thus 2 and 3 only is the correct answer.

First electron affinity and second electron affinity are definitions that need to be learned:

First electron affinity is the enthalpy change when one electron is gained by each atom in one mole of gaseous atoms.

Second electron affinity is the enthalpy change when one electron is gained by each 1− ion in one mole of gaseous ions.

Thus Cl_(g) + e⁻ → Cl⁻_(g) is the correct answer, as this is the first electron affinity of chlorine.

Across any period there is a general decrease in atomic radii due to increasing nuclear charge and electrons filling the same energy level. The electrons are in the same energy level, but the nuclear charge increases (number of protons increases), increasing the attraction between the nucleus and the electrons and pulling them in more tightly/closer to the nucleus. Overall the attraction between the outer electrons and the nucleus outweighs the increasing electron-electron repulsion.

Down any group there is a general increase in atomic radii due to electrons filling energy levels that are further from the nucleus making the atoms bigger, despite the increase in nuclear charge.

Thus K > Na > Al > Si is the correct answer. K is lower than sodium in group 1, so Na will have a smaller atom radius, and Al and Si are elements further across period 3 than Na, so will have decreasing atomic radii.

Metal oxides (CaO), if they react with water at all, will give alkaline solutions (this does react).

Non-metal oxides (P₄O₁₀ and SO₃), if they react with water at all, will give acidic solutions (both of these do react).

3 only is therefore the correct answer.

Group 2 are metals, so have metallic bonding and therefore giant structures. The melting and boiling points become lower down the group because of the increasing atomic/ion size leads to less attraction.

Group 18 are non-metals; the noble gases, and exist as atomic gases. The melting and boiling points become higher down the group as London dispersion forces get stronger as the number of electrons/molar mass of the atoms increases.

Electronegativities decrease down all groups (and increase across periods).

Ionisation energies decrease down all groups (and show a general increase across periods).

Thus Ionisation energies decrease is the correct answer.

Metals and Non-metals (or metalloids) with giant structures (Li, Be, B and C) have high melting/boiling points; strong covalent or metallic bonds are broken. Non-metals with simple molecular structures (N₂, O₂, F₂) or gaseous atoms (Ne) have low melting/boiling points; only intermolecular forces need to be broken. So statement 2 is false as boiling point will peak at C (in the middle) across period 2.

Across any period there is a general decrease in atomic radii due to increasing nuclear charge and electrons filling the same energy level. The electrons are in the same energy level, but the nuclear charge increases (number of protons increases), increasing the attraction between the nucleus and the electrons and pulling them in more tightly/closer to the nucleus. Overall the attraction between the outer electrons and the nucleus outweighs the increasing electron-electron repulsion. Statement 1 is correct.

Across any period there is a general increase in 1^st ionization energy due to increasing nuclear charge and electrons filling the same energy level. So the outermost electron becomes subject to greater attraction and more difficult to remove. Statement 3 is correct.

Thus 1 and 3 only is the correct answer.

Metalloids lay on the diagonal between metals and non-metals. Beginning with Boron each successive element on the diagonal, and the one underneath it, are considered to be metalloids, ending in Astatine (which is not considered to be a metalloid):

Group 1 metals react exothermically with water to give strongly alkaline solutions of the metal hydroxide and hydrogen gas, e.g. 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

Therefore 'The reaction is exothermic and the products are sodium hydroxide and hydrogen' is the correct answer.

The most electronegative element is Fluorine (4.0) and the least is Francium (0.7). Therefore, electronegativity increases across all periods and up all groups. You don't need to know the values, but you do need to know the trend.

Therefore the correct answer is 'Electronegativity increases across period 2 and decreases down group 17'.