Inorganic Chem

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General advice for students in inorganic chemistry:

Ensure that a good understanding of the foundational topics is present before attempting this section, for this is a pre-requisite to scoring well. The answers required tend to be qualitative, and the student should pay particular attention to use of key concepts and key words in structuring his answer. Model answers found in past year prelim papers provide a good guide to this.

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Group 2 – FAQ

Q: Would you expect Group II elements to react more vigorously with water than the Group I elements
from the same period?
A: In order to answer this question, one has to compare the standard reduction potential of both the
Groups I and II elements. The more negative the standard reduction potential, the less feasible the
reduction of the cation and hence, the more feasible for the aqueous cation to form from the metal.
Hence, the reaction between water and Group II metal should be more vigorous.

Q: Why do the decomposition reactions of Group II carbonate, nitrate and hydroxide all result in the
formation of the same product, the oxide?
A: It goes to show that the oxide formed is relatively more stable than the corresponding Group II
carbonate, nitrate or hydroxide compounds. This is the driving force for decomposition — the formation
of a more stable compound. In addition, the production of gaseous products increases the degree of
disorderliness, hence the entropy of the system.

Q: How can we know that the oxide is more stable than the other three types of compounds? How is
stability measured?
A: Since Group II compounds (with the exception of Be compounds) are ionic, the stability of the solid
compound is dependent on the strength of the ionic bonds, which is indicated by the magnitude of the

lattice energy. In general, the more exothermic the lattice energy, the stronger the ionic bonds, the more
stable the compound, the more difficult it is to decompose.
Lattice energy is proportional to (+q)(-q)/(r+ + r-)

where q is charge and r is ionic radius.
Based on the above equation, it goes to show that an ionic compound with a higher ionic charge and a
smaller ionic radius will tend to be more stable than a compound with smaller ionic charge and larger
ionic radius.
The underlying reason is that the greater the charges on the ions, the stronger the attraction between the
oppositely charged ions will be, and stronger attraction gives rise to a more stable ionic compound. With
smaller sizes, oppositely charged ions can be closer to each other in the solid lattice, resulting in stronger
attractive forces and hence, leads to a more stable compound.

Q: BeO and Be(OH)2 are amphoteric. We can use the reaction with NaOH(aq) to test for their acidic
property. How do we test for their basic property?
A: Well, one just needs to add a mineral acid to it. If it is a base, one would get salt and water being
formed.

Q: Would you expect beryllium salts to be more thermally unstable or less thermally unstable compared
to those of the other Group II elements?
A: Beryllium salts are more thermally unstable due to the higher charge density of the Be2+, which gives
rise to greater polarizing power. This would cause the electron cloud of the anion that is attracted to the
Be2+ ion to be more distorted, causing the covalent bonds within the anion to be weakened to a greater
extent. Thus, beryllium salts are easier to decompose.

Periodicity – FAQ

Q:  Explain the following observations using equations whenever possible:
(i) Carbon dioxide gas is produced on adding sodium carbonate to aluminum chloride solution, but not on adding sodium carbonate to sodium chloride solution.

NaCl dissolves in water to give neutral solution. Only hydration takes place. Due to the low charge density of Na+, it will not undergo hydrolysis in water.

AlCl3 dissolves in water with slight hydrolysis to give acidic solution.
Acidic nature of AlCl3 due to highly charged Al3+ which polarizes H2O molecules to liberate H+ ions.

AlCl3(s) + 6H2O  [Al(H2O)6]3+(aq) + 3Cl–(aq)
[Al(H2O)6]3+(aq)  [Al(H2O)5(OH)]2+ + H+

CO2 gas produced when Na2CO3 added to acidic AlCl3 solution.
Na2CO3 + 2H+  2Na+ + H2O + CO2

(ii) An acidic solution is produced on adding water to silicon tetrachloride, whereas an oily layer is observed on adding water to carbon tetrachloride.

SiCl4 undergoes complete hydrolysis in water to give acidic solution.
SiCl4 + 2H2O  SiO2 + 4HCl
CCl4 does undergo hydrolysis with water. Hence, it is insoluble in water.

(b)(i) The element aluminium and its compounds have some properties characteristic of metals, and some of non-metals. Aluminium hydroxide, for example, is known to be amphoteric. Explain what the word in italics means.

An amphoteric compound is one which reacts with both acids and bases.

(ii) Aluminium sulphate and calcium oxide are sometimes added to water supplies to
precipitate suspended solids and bacteria. A small amount of aluminium-containing
ions remains in the solution and its presence in drinking water may contribute to the
mental illness known as Alzheimer’s disease.
Write a balanced equation for the reaction that occurs when aluminium sulphate and
calcium oxide is added to water, given that aluminium hydroxide is one of the products
formed. 

Al2(SO4)3 (s) + 3CaO (s) + 3H2O (l)® 2Al(OH)3 (s) + 3CaSO4 (s) 

(iii) Explain why adding too much calcium oxide would increase the probability of contracting Alzheimer’s disease. Write equations for all reactions that occur.

CaO dissolves only slightly to form an alkaline solution of Ca(OH)2.
CaO + H2O ® Ca(OH)2
Al(OH)3 (s) + OH-  Al(OH4)-
Q: State and explain how each of the following properties varies across the third period of
the Periodic Table from Na to Cl :
(i) Ionic radius for the elements

Cations are always smaller than their parent atoms. E.g. atomic radius of Na > Ionic radius of Na+.

Number of protons is the same as that of the neutral atom.

. Hence, nuclear charge is constant
· Valence electrons are removed from the atoms. The cations formed have one less shell than the neutral atoms.
· The remaining outer electrons are closer to the nucleus and are more strongly attracted by the nucleus.
· This causes the ionic radii of the positive ions to be smaller than the corresponding atomic radii of its neutral atoms.
But, anions are always larger than their parent atoms. E.g. atomic radius of Cl < Inoic radius of Cl-
· The number of protons in the ion remains the same as in the neutral atom, hence nuclear charge remains constant.
· The anions have more electrons than protons so the electrostatic forces of attraction on the outer electrons is less than that in neutral atoms
· Hence, the outer electrons are less strongly attracted by the nucleus.
· Thus, the ionic radii of negative ions are larger than the corresponding atomic radii of its neutral atoms.

(ii) boiling point of chlorides

· The bonding of chlorides changes from ionic to covalent across the period, because the electronegativity difference between the element and chlorine decreases.
· NaCl and MgCl2 have giant ionic structures.

Hence they have high boiling point. As large amount of energy are required to overcome the extensive and strong electrostatic forces of attraction between oppositely charged ions present in the crystal lattice.
· Al2Cl6 is a covalent compound. It has simple molecular structure.

Hence, it has a low boiling point as small amount of energy is required to overcome the weak van der Waals’ forces of attraction between molecules during melting.

· SiCl4, PCl3 and PCl5 are covalent compounds.
· They have simple molecular structures.

These chlorides have lower boiling points as small amount of energy is required to overcome the weak van der Waals’ forces of attraction between the molecules during melting.

(iii) electronegativity

· Electronegativity increases across the period from Na to Cl as effective nuclear charge increases and atomic radius decreases across the period.
(iv) electrical conductivity of elements

Na, Mg and Al are giant metallic structures. They are good conductors because of the mobile sea of delocalised electrons. Current is carried by these delocalised valence electrons which function as charge carriers when a potential difference is applied. Conductivity increases as more valence electrons are added to the sea of electrons from Na to Mg to Al.
Si has a giant molecular structure with low electrical conductivity.
P, S, and Cl have simple molecular structures. P, S and Cl are non-metals and non-conductors of electricity. There is an absence of mobile electrons in these elements as their valence electrons are involved in covalent bonding and are fixed in position.

(b) “Beryllium differ from other group II metals but shows a strong resemblance to aluminium instead.”

Comment on this statement by referring to the bonding of oxides and chlorides of Beryllium. 
There is a diagonal relationship between Be and Al because they have similar electronegativity and charge density. Be resembles Al in some of their reactions where both BeCl2 and AlCl3 are covalent compounds and BeO and Al2O3 are amphoteric in nature where they form amphoteric oxides.