Sem1 Unit5 Chemical Bonding

SK017 Unit 5

Past Year Examination Questions

Unit 5: Chemical Bonding Jan 99 1.

With reference to the structure/lattice and bonding, discuss the electric conductivity of: i. magnesium ii. graphite iii. sodium chloride

2.

For SF4, NH3 and CCl4 molecules, draw their shapes and state whether these molecules are polar or non-polar.

3.

Elements phosphorus and nitrogen are in the group 15 of the Periodic Table. i. Draw the Lewis structures for PCl5 and NCl3 molecules. ii. Explain why phosphorus can form compounds PCl3 and PCl5 but nitrogen can only form NCl3.

4.

Elements combine among themselves or with other elements to attain their stabilities. By using fluorine and magnesium elements as examples, discuss how the following bondings are formed: i. electrovalence / ionic bond ii. covalent bond iii. metallic bond

5.

Fluorine and chlorine are elements from the group 17 while boron is an element from the group13. Boron reacts with fluorine to form BF3 compound. Chlorine reacts with fluorine to form ClF3 compound. i. Compare the bond angles for FBF in BF3 with FClF in ClF3. Explain. ii. What are the hybridisations of B atom in BF3 and Cl atom in ClF3.

Jan 00 6.

How would an atom acquire an octet arrangement when it forms bond with the other atom?

7.

Xenon, Xe in the noble gas group has octet electronic configuration. Xe is able to form compounds like fluorides with formulae XeF2 and XeF4. a) State the number of bonding electron pairs and lone electron pairs which surround the central atom Xe in XeF2 and XeF4. Why Xe atom allows these numbers of electrons surrounding it? b) Give the shapes of XeF2 and XeF4 molecules. c) State the hybridisation of Xe atom in XeF2 and XeF4. d) Apart from XeF2 and XeF4, give another molecular formula for a fluoride compound of Xe.

8.

What is meant by hybridisation? Sketch and label the diagram of orbital overlapping in ethene, C2H4 and explain the hybridisation of the carbon atom.

9.

State the factors that determine the polarity of molecules. State and explain which compound has the most ionic character amongst the compounds: N2O4, H2O2, HF, CO2 and IBr.

25

SK017 Unit 5

Past Year Examination Questions

10.

The table below shows two elements with their respective proton numbers. Element Proton number

U 20

V 9

Based on the above table, what types of bonds are formed between elements U and V. Explain how the bonding is formed. 11.

For nitrate ion, NO3−, draw the resonance structures; give the shapes of the ion and state the hybridisation of the central atom. Determine the polarity of this ion and indicate the direction of the dipole moment on its structural formula.

12.

Arrange in order of increasing boiling points for the solid crystals of copper, Cu, iodine, I2 and diamond, C. Explain your answer with reference to the attractive forces between atoms. Which of the solid crystals can conduct electricity in the solid state? Explain.

13.

a) For SF4 molecule i. Draw its Lewis structure ii. State the number of electron pairs on the central atom. iii. Predict the actual geometry of the molecule. iv. What is the hybridisation of the central atom S? b) Iodine can form compounds with more than one oxidation state like in ICl3 and ICl4− molecule. i. Draw the shape of ICl4− ion. ii. Explain the difference in bond angles between ICl3 and ICl4−.

14.

Boron hydride, BH3 and ammonia, NH3 react in the ratio 1:1 to form a product. a. Explain how the bond can be formed between NH3 and BH3 by means of Lewis structures. b. What type of bonding is formed between BH3 and NH3? c. What happen to the bond angles in ammonia and BH3 after forming the product? Explain. d. Compare the boiling point of BH3 with NH3. Explain the differences. e. What types of hybridisation are experienced by the element B in BH3 and N in NH3?

Jun 00 15.

Given the Lewis structures of three resonance structures of thiocyanate ion, SCN−: S

C

N

S

C

N

S

C

N

I II III i. Calculate the formal charge of each atom in each resonance structure. ii. Give the most plausible Lewis structure. 16.

Given the electronic configurations for hydrogen atom and phosphorus atom: H : 1s1 P : 1s2 2s2 2p6 3s2 3p3 i. Sketch the molecular geometry of phosphine, PH3 and state its shape. ii. State whether the PH3 molecule is polar.

17.

Determine which of the compounds, hydrogen bromide, HBr or carbon tetrafluoride, CF4 has a higher boiling point. Explain. 26

SK017 Unit 5

Past Year Examination Questions

18.

How the covalent bond and the electrovalence bond are formed? State the factors that affect the strength of electrovalence bond. Discuss the strength of electrovalence bond for Na2O and MgO.

19.

The central atom of IF4+ ion does not obey the octet rule. Show the orbital diagram for the central atom in the ground state and excited state. Show the hybridisation that occurs for the central atom of this ion as well as the overlapping of its orbitals with the orbitals of the terminal atoms. By using the valence-shell electron-pair repulsion (VSEPR) theory, show how the geometry for this species can be determined. Sketch its geometry.

20.

Show how the Lewis structures for phosphate ion, PO43− can be obtained. Sketch and name the geometry of phosphate ion. Hence compare the geometry of phosphate ion with that of sulphite ion, SO32−. Draw the Lewis structure for magnesium phosphate compound. Identify the types of bonding that exist in the compound. Explain.

Mac 01 21.

An element of iodine combines with chlorine to form ICl2+ and ICl2− ions. i. Draw the Lewis structures and predict the shapes for these two ions based on the valence-shell electron pair repulsion (VSEPR) theory. ii. State the hybrid orbitals for the central atom iodine, I in both ions. iii. Determine whether these two ions are polar.

22.

Explain the van der Waals forces and give its types. Explain how the atomic size or molecular size can influence the strength of these forces.

23.

By drawing suitable diagrams, explain all the possible types of hybrid orbitals formed between s orbital and p orbitals. Draw an orbital diagram for benzene molecule, and label all the σ- and π-bonds.

Aug 02 24.

The table below shows the melting points, boiling points and solubility in water and CCl4 for compounds E and F. Compoun d E F i. ii. iii.

25.

Melting point (°C) 2800 -25

Boiling point (°C) 3600 -10

Solubility in water

Solubility in CCl4

High Low

Very low High

Predict the type of compound for E and F. The boiling point of compound E is higher than that of compound F. Explain. Compound E has a higher solubility in water. Explain.

Give the factors that influence the strength of the van der Waals’ forces. Oxygen and sulphur are elements in the group 16 in the Periodic Table. Compare the boiling points of H2O and H2S based on the intermolecular forces.

27

SK017 Unit 5

Past Year Examination Questions

26.

The proton numbers for elements S, T and U are 9, 17 and 38 respectively. Write the electronic configurations for element T and U. Give the formula for the compound formed when element T and U combine as well as state the type of bonding formed. If TS4− and TS2+ ions exist, Draw the Lewis structures for these two ions and determine the hybridisations of T atom in TS4− and TS2+. Hence predict the possible shape for both ions.

Sept 03 27.

Phosgene, COCl2 is a colourless and highly toxic gas. Draw all the possible Lewis structures for COCl2 and determine the most likely (plausible) resonance structure by showing the formal charges on each atom.

28.

(a) Explain the hybridisation of phosphorus atom in phosphorus pentachloride, PCl5. Hence, show the shape of the molecule. (b) Sodium chloride, metallic copper, diamond and sulphur are four examples of solid substances which have a giant structure. With the aid of diagrams, show the bonds that built the giant structure for each of the substances.

29.

(a) Using suitable examples, explain the formation of compounds through covalent and dative covalent bondings. (b) By using a suitable example, explain hydrogen bonding and its effects on two physical properties of the substance.

Oct 04 30.

Consider the SF4 molecule. Draw the orbital energy level diagram for valence shell electrons of sulphur in the ground state and the hybridised state.

31.

What is meant by resonance hybrid? i. Write two resonance structures for NCO− ion. Designate formal charge on each atom, if any. ii. Which structure is more stable? Explain.

32.

The following table apply to the compounds QClx and RCly. Compound QClx RCly

Melting point (°C) 605 -25.2

Boiling point (°C) 1343 86.0

Solubility in water

Solubity in CCl4

38 0.07

0.071 Complete miscible

Compare the physical properties of QClx and RCly in terms of bonding, physical state at room temperature, volatility and solubility. Explain. Oct 05 33.

Nitrogen and phosphorus are in group 15 of the Periodic Table and can form covalent compounds. i. What is the maximum number of covalent bonds that can be formed by a central atom of nitrogen and phosphorus respectively? Explain the differences. ii. Draw the Lewis structures for NH3 and PCl5.

28

SK017 Unit 5

Past Year Examination Questions

34.

The reaction of xenon, Xe with fluorine, F forms xenon tetrafluoride, XeF4. i. Draw the structure of XeF4. ii. Predict the polarity of the bonds and the molecule of XeF4. Explain.

35.

(a) Explain the formation of bonding in HCOOH molecule using the hybridisation and orbital overlapping concepts. (b) Explain the differences between covalent bonding and dative covalent bonding using AlCl3 and Al2Cl6 molecules as examples.

Oct 06 36.

(a) Explain why the boiling point of Cl2 (-34.6˚C) is lower than that of Br2 (58.8˚C) and the vapour pressure of acetone, CH3COCH3 is higher than that of ethanol, CH3CH2OH at room temperature. (b) There are two possible structures for BeCl2 but the more stable structure is one which does not obey the octet rule. Explain your answer.

37.

(a) What is meant by hybridisation? Cyanide ion, CN- is formed when the carbon atom undergoes hybridisation and bonds to nitrogen atom. Show the hybridisation in CN- and draw the overlapping of orbitals. (b) Briefly describe metallic bonding. The boiling point of sodium, magnesium and aluminium are 881, 1105 and 2467˚C respectively. Explain the differences in the boiling points.

Oct 07 38.

(a) Draw the possible Lewis structures for COCl2 and determine the more stable structure. Explain your answer. (b) Predict the shape of IF4+ ion.

39.

What is meant by hybridisation? Glycine, NH2CH2COOH is one of the essential amino acids. State the types of hybridisation of C, N and O atoms in the glycine molecule. Draw and label the overlapping of orbitals showing all the σ and Π bonds formed in the glycine molecule. Predict the C-N-H and C-C-O angles in the glycine molecule. Explain your answer.

Oct 08 40.

(a)

Explain why ammonia, NH3, has a higher boiling point than methane, CH4.

(b)

i. Draw the orbital diagram for phosphorus in PCl6-. ii. Name the hybrid orbital used by phosphorus in PCl6-. iii. Draw and name the geometry of PCl6- ion.

29

SK017 Unit 5

Past Year Examination Questions

41.

The proton number of element J and K are 13 and 16 respectively. Draw the orbital diagram for the valence electrons of each element. Suggest the most stable ions for J and K. Write their respective electronic configurations.

42.

(a) Magnesium is a good electrical conductor. Draw a diagram of an electron sea model to explain the bonding formed in the metal. How would the model explain the electrical conductivity of the metal? (b) Based on the skeletal structure of peroxynitrite ion, OONO-, draw all the possible Lewis structures. Assign formal charge to each atom in the structure. Determine the most stable Lewis structure for the ion and explain your answer. What is the hybridisation of the N atom in the peroxynitrite ion? Estimate the O-O-N and O-N-O bond angles.

30