Month: April 2020

Question 1

Aluminium and thallium are elements in Group 3 of the Periodic Table. Both elements form compounds and ions containing chlorine and bromine.

1. Write an equation for the formation of aluminium chloride from its elements.
2. An aluminium chloride molecule reacts with a chloride ion to form the AlCl₄⁻ ion. Name the type of bond formed in this reaction. Explain how this type of bond is formed in the AlCl₄⁻ ion.
3. Aluminium chloride has a relative molecular mass of 267 in the gas phase. Deduce the formula of the aluminium compound that has a relative molecular mass of 267
4. Deduce the name or formula of a compound that has the same number of atoms, the same number of electrons and the same shape as the AlCl₄⁻ ion.
5. Draw and name the shape of the TlBr₅²⁻ ion.
6. Draw the shape of the TlCl₂⁺ ion.
7. Explain why the TlCl₂⁺ ion has the shape that you have drawn in part (f).

Question 2

Ammonia gas readily condenses to form a liquid when cooled.

1. Name the strongest attractive force between two ammonia molecules.
2. Draw a diagram to show how two ammonia molecules interact with each other in the liquid phase. Include all partial charges and all lone pairs of electrons in your diagram.
3. Ammonia reacts with boron trichloride to form ammonia boron trichloride. Draw the molecule and state how the bond between ammonia and boron trichloride is formed.
4. The following table shows the electronegativity values of some elements.

		H	Li	B	C	O	F
	Electronegativity	2.1	1.0	2.0	2.5	3.5	4.0

 

1. Give the meaning of the term electronegativity.
2. Suggest the formula of an ionic compound that is formed by the chemical combination of two different elements from the table.
3. Suggest the formula of the compound that has the least polar bond and is formed by chemical combination of two of the elements from the table.

Question 3

Hydrogen peroxide is a very good oxidising agent.

1. Suggest a value for the H−O−O bond angle.
2. Hydrogen peroxide dissolves in water.

1. State the strongest type of interaction that occurs between molecules of hydrogen peroxide and water.
2. Draw a diagram to show how one molecule of hydrogen peroxide interacts with one molecule of water. Include all lone pairs and partial charges in your diagram.

3. Explain, in terms of electronegativity, why the boiling point of H₂S₂ is lower than H₂O₂.

Question 4

The following table shows the electronegativity values of the elements from lithium to fluorine.

		Li	Be	B	C	N	O	F
	Electronegativity	1.0	1.5	2.0	2.5	3.0	3.5	4.0

1. State the meaning of the term electronegativity.
2. Suggest why the electronegativity of the elements increases from lithium to fluorine.
3. State the type of bonding in lithium fluoride. Explain why a lot of energy is needed to melt a sample of solid lithium fluoride.
4. Deduce why the bonding in nitrogen oxide is covalent rather than ionic.
5. Oxygen forms several different compounds with fluorine.

1. Suggest the type of crystal shown by OF₂
2. Write an equation to show how OF₂ reacts with steam to form oxygen and hydrogen fluoride.
3. One of these compounds of oxygen and fluorine has a relative molecular mass of 70.0 and contains 54.3% by mass of fluorine. Calculate the empirical formula and the molecular formula of this compound. Show your working.

Question 5

Fritz Haber, a German chemist, first manufactured ammonia in 1909. Ammonia is very soluble in water.

1. State the strongest type of intermolecular force between one molecule of ammonia and one molecule of water.
2. Draw a diagram to show how one molecule of ammonia is attracted to one molecule of water. Include all partial charges and all lone pairs of electrons in your diagram.
3. Phosphine (PH₃) has a structure similar to ammonia. In terms of intermolecular forces, suggest the main reason why phosphine is almost insoluble in water.

Question 1

Aluminium and thallium are elements in Group 3 of the Periodic Table. Both elements form compounds and ions containing chlorine and bromine.

1. Write an equation for the formation of aluminium chloride from its elements.
2. An aluminium chloride molecule reacts with a chloride ion to form the AlCl₄⁻ ion. Name the type of bond formed in this reaction. Explain how this type of bond is formed in the AlCl₄⁻ ion.
3. Aluminium chloride has a relative molecular mass of 267 in the gas phase. Deduce the formula of the aluminium compound that has a relative molecular mass of 267
4. Deduce the name or formula of a compound that has the same number of atoms, the same number of electrons and the same shape as the AlCl₄⁻ ion.
5. Draw and name the shape of the TlBr₅²⁻ ion.
6. Draw the shape of the TlCl₂⁺ ion.
7. Explain why the TlCl₂⁺ ion has the shape that you have drawn in part (f).

Question 2

Ammonia gas readily condenses to form a liquid when cooled.

1. Name the strongest attractive force between two ammonia molecules.
2. Draw a diagram to show how two ammonia molecules interact with each other in the liquid phase. Include all partial charges and all lone pairs of electrons in your diagram.
3. Ammonia reacts with boron trichloride to form ammonia boron trichloride. Draw the molecule and state how the bond between ammonia and boron trichloride is formed.
4. The following table shows the electronegativity values of some elements.

		H	Li	B	C	O	F
	Electronegativity	2.1	1.0	2.0	2.5	3.5	4.0

 

1. Give the meaning of the term electronegativity.
2. Suggest the formula of an ionic compound that is formed by the chemical combination of two different elements from the table.
3. Suggest the formula of the compound that has the least polar bond and is formed by chemical combination of two of the elements from the table.

Question 3

Hydrogen peroxide is a very good oxidising agent.

1. Suggest a value for the H−O−O bond angle.
2. Hydrogen peroxide dissolves in water.

1. State the strongest type of interaction that occurs between molecules of hydrogen peroxide and water.
2. Draw a diagram to show how one molecule of hydrogen peroxide interacts with one molecule of water. Include all lone pairs and partial charges in your diagram.

3. Explain, in terms of electronegativity, why the boiling point of H₂S₂ is lower than H₂O₂.

Question 4

The following table shows the electronegativity values of the elements from lithium to fluorine.

		Li	Be	B	C	N	O	F
	Electronegativity	1.0	1.5	2.0	2.5	3.0	3.5	4.0

1. State the meaning of the term electronegativity.
2. Suggest why the electronegativity of the elements increases from lithium to fluorine.
3. State the type of bonding in lithium fluoride. Explain why a lot of energy is needed to melt a sample of solid lithium fluoride.
4. Deduce why the bonding in nitrogen oxide is covalent rather than ionic.
5. Oxygen forms several different compounds with fluorine.

1. Suggest the type of crystal shown by OF₂
2. Write an equation to show how OF₂ reacts with steam to form oxygen and hydrogen fluoride.
3. One of these compounds of oxygen and fluorine has a relative molecular mass of 70.0 and contains 54.3% by mass of fluorine. Calculate the empirical formula and the molecular formula of this compound. Show your working.

Question 5

Fritz Haber, a German chemist, first manufactured ammonia in 1909. Ammonia is very soluble in water.

1. State the strongest type of intermolecular force between one molecule of ammonia and one molecule of water.
2. Draw a diagram to show how one molecule of ammonia is attracted to one molecule of water. Include all partial charges and all lone pairs of electrons in your diagram.
3. Phosphine (PH₃) has a structure similar to ammonia. In terms of intermolecular forces, suggest the main reason why phosphine is almost insoluble in water.

1. Explain why helium’s first ionisation energy is much higher than hydrogen’s.
2. Explain why lithium has lower ionisation energy than helium even though lithium has 1 proton more.
3. Why does the main trend of first ionisation energy increase across a period?
4. Explain why the ionisation energy of aluminium is lower than that of magnesium.
5. Explain why the first ionisation of magnesium is 736 kJ mol^-1 whilst that of beryllium is 900 kJ mol^-1.
6. Explain why the second ionisation energy of sodium is much higher than the second ionisation energy of magnesium.
7. The four successive ionisation energies of beryllium are: 900 1760 1480 21000 (kJ mol^-1). Explain why there is a very big increase in ionisation energies between the second and third ionisation energies.
8. Give the group number for the following elements having these successive first ionisation energies.
   1. 799 2420 3660 25040 32850
   2. 1000 2260 3390 4540 6990 8790 27100 31700
   3. 736 1450 3740 7740 10500 1800 21700 25600
9. Which would have a greater value for the second ionisation energy: magnesium or aluminium?
10. Which would have a greater value for the third ionisation energy: magnesium or aluminium?

 

Atomic Structure

Question 1

Write the electronic configuration for: C, N^3-, Fe²⁺, Cu, Cu⁺, I and Pb.

Question 2. 

Put the following atoms/ions in order of decreasing size. Explain your answer. The atoms/ions are: Na⁺, Al³⁺, Ne, Ar, K⁺, N^3-.

Question 3:

Draw a graph representing all the ionisation energies possible for the Al atom.

Bonding

Question 1

Explain the trends in boiling points for the hydrogen halides.

Question 2

Draw the following molecules/ions: BH₂^–, NI₃, ClF₄, SF₅^–, BO₃^3-, O₃²⁺, IO₂⁺, BBr₃^2-

Energetics

Question 1

On strong heating, calcium carbonate decomposes to calcium oxide and carbon dioxide:

Owing to the conditions under which the reaction occurs, it is not possible to measure the enthalpy change directly.

An indirect method employs the enthalpy changes when calcium carbonate and calcium oxide are neutralized with hydrochloric acid.

i) Write the equation for the reaction of calcium carbonate with hydrochloric acid. State symbols are not required. [ΔH1 is the enthalpy change for this reaction]

ii) The reaction of calcium oxide with hydrochloric acid is

Use the equations in parts (i) and (ii) to complete the Hess’s Law cycle below to show how you could calculate the enthalpy change for the decomposition of CaCO₃. Label the arrows in your cycle.

Question 2

Calculate ΔH and ΔS  for the following reaction:

Use the results of this calculation to determine the value of ΔG^o for this reaction at 25^o C, and explain why NH₄NO₃ spontaneously dissolves in water at room temperature.

Compound              H_f^o(kJ mol^-1)         S(J mol^-1 K^-1)

NH₄NO₃(s)                 -365.56              151.08

NH₄⁺ (aq)                  -132.51               113.4

NO₃^– (aq)                   -205.0                 146.4

 

Equilibria

Kc

Question 1

Esters are a useful group of compounds due to their distinctive smells. One example
of an ester is ethyl ethanoate, its formation is shown below.

a) Systems like this are described as being a ‘dynamic equilibrium’. Explain
the term ‘dynamic equilibrium’

b) Write down the expression for the equilibrium constant, Kc, for this reaction.

c) Calculate the value of Kc for this reaction given the equilibrium concentrations below.

[CH₃COOH] = 0.08 moldm^-3, [C₂H₅OH] = 0.08 moldm^-3, [CH₃COO C₂H₅] = 0.25 moldm^-3, [H₂O] = 0.1 moldm^-3

d) Concentrated sulphuric acid is added to the reaction mixture as it removes water molecules. What effect would this have on the equilibrium position of this system?

Question 2

A mixture of 1.441 g of H₂ and 70.24 g of Br₂ is heated in a 2.00-L vessel at 700 K. These substances react as follows.

At equilibrium, the vessel is found to contain 0.627g of H₂. Calculate their equilibrium concentrations and Kc.

Kp

Question 1

Ammonium iodide dissociates reversibly to ammonia and hydrogen iodide.

At 400oC,Kp=0.215. Calculate the partial pressure of ammonia at equilibrium when a sufficient quantity of ammonium iodide is heated to 400oC.

Question 2

For the reaction

at 550 K, the equilibrium constant (Kp) is 9.81kPa. Suppose that 3.150 g AsCl₅ is placed in an evacuated 600 ml bulb, which is then heated to 550K.

What is the partial pressure of AsCl₅ at equilibrium? (For this question you would need to use Pv=nRT and in the ICE table, the pressure is used instead of moles)