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ST. JOSEPH SEMINARY – SANU

FORM FIVE HOME PACKAGE – 2020

CHEMISTRY

1. (a) Explain the concept of Dalton atomic theory.

(b) Establish evidence for electronic structure of an atom.

(c) Describe atomic model according to:

(i) Rutherford

(ii) Bohr

(iii)J. J. Thomson

2. (a) What is isotopes?

(b) Discuss significance of isotopes in real life situation

(c) Describe energy levels and hydrogen spectrum

(d) Work out characteristics of atomic spectrum

3. (a) Explain subatomic particles

(b) Compare mass number with atomic number

(c) Describe mass spectroscopy

4. (a) Describe the wave particle duality of the electron

(b) Derive the debroglie equation

(c) Explain quantum numbers:

(i) n (ii) l, (iii) ml (iv) Ms

5. (a) Describe atomic orbitals and their respective shapes

(b) Explain the rules responsible for filling electrons in a ground state

6. (a) What is meant by chemical bonding?

(b) Differentiate the characteristics of electrovalent bond and covalent bond

(c) Describe dative bonding

(d) Explain the meaning and significance of hydrogen bonding

(e) Explain briefly the factors that control the formation of:\

(i) Covalent bond

(ii) Electrovalent bond

7. (a) What is meant by hybridization of atomic orbital

(b) What are the characteristics of hybrid orbital

(c) Describe the shapes of:

(i)

(ii) orbital

(d) Explain the overlapping of hybrid orbital to form a molecular orbital

8. (a) Explain how mass spectrometer works.

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(b) What is meant by:

(i) Stationary state

(ii) Excited state

(iii)Hyperfine live

(iv) Zeeman effect

(v) Stark effect

(vi) Heisenberg uncertainty principle with its related Mathematical expression.

(c) Give application of quantum number.

9. (a) Magnesium has three stable isotopes masses 23.985, 24.985, 25.982 a.m.u. The relative

abundances of the three isotopes are 39.35, 5.065 and 5.585 respectively. Calculate the

average atomic mass.

(b) A sample of pure unknown element “X” was analysed and the data is given in the table

below. Calculate the relative atomic mass of “X”. Strictly to two decimal places.

Symbol Mass (a.m.u) Natural abundance (%)

19.992 90.22

20.994 0.257

21.991 8.82

10. (a) The mass spectrum of an element enables the relative abundance of each isotope of the

element to be determined. Data relating to the mass spectrum of element “X” whose

atomic number is 35 appear as follows:

Mass number of isotopes Relative Abundance

79 50.5%

81 49.5%

(i) Write down the conventional symbols for the two isotopes of “X”.

(ii) Calculate the relative atomic of “X” to the three significant figures.

(b) Silver consist of two isotopes

of atomic masses 106.91g/mole and

108.91g/mol respectively. The relative abundances of these isotopes are 51.88% for

and 48.12% for

Calculate the average atomic mass of Ag.

11. (a) Copper has two naturally occurring isotopes Cu-63 have atomic mass of 62.9296 a.m.u

and an abundance of 69.15%. What is the atomic mass of the second isotope? Give that

average atomic mass of Copper is 63.546 a.m.u

(b) In a sample of 400 lithium atoms. It is found that 30 atoms are lithium -6(6.015g/mol)

and 370 atoms are Lithium -7(7.016g/mol). Calculate the average atomic mass of

Lithium).

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12. (a) The diagram below shows the mass spectrum of Magnesium. The heights of the three

peaks and the mass number of isotopes are shown. Calculate the relative atomic mass of

Magnesium

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Intensity

(Isotopic %) 9.1

8.1

24 25 26 Mass number

(b) Natural rubidium has the average mass 85.4678 a.m.u and compared of isotope

( ) . The ratio of atoms

in

natural rubidium is 2.591.

13. (a) What is electromagnetic spectrum?

(b) Discuss the relationship between wavelength, frequency and speed of radiation.

14. (a) Most of the light from a Sodium vapour lamp has wavelength of 589nm. What is:

(i) Frequency and

(ii) Wave number of this radiation (b) Explain the differences between Emission spectrum and Absorption spectrum.

15. (a) Calculate the wavelength of the first line in Balmer series. ).

(b) Calculate the energy of a photon of radiations whose wavelength is from

sodium atoms when heated.

Given that,

(c) Using Plank’s equation, calculate the energy of photons of light for radiation of wave

length 242.4nm, the longest wavelength that will bring about the photo dissociation of

What is the energy of:

(i) One photon

(ii) A mole of photons of this light?

(

16. (a) If the wavelength of the first member of Balmer series in hydrogen spectra is Calculate the wavelength of the first member of the Lyman series.

(b) (i) Give differences between line and band spectra.

(ii) Explain where the different colours of light come from in the bright linen spectrum of

an element.

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17. (a) A hydrogen atom with an electron in its ground state interacts with a photon of light with

a wavelength of . Could the electron make a transition from the ground state

to a higher energy level? If it does make a transition, indicate which one. If no transition

can occur, explain.

(b) What is the wavelength (in mm) of the radiation required for a hydrogen electron

transition from the 1s orbital to a 2p orbital?

(c) (i) Calculate the frequency and the wavelength of a spectra line of hydrogen

corresponding to a transition of electron from (ii) In what region of if the elemagnetic spectrum would this spectral line be?

18. (a) A diode laser emits at wavelength of 987nm. All the radiation it emits is absorbed in

detector which measures a total energy of over a period of 32 seconds. How many

photons per second are being emitted by the laser?

(b) Calculate the wavelength of the line in Balmer series associated with energy transitions.

where,

19. (a) An electromagnetic radiation was emitted in Balmer series as result of electron transition

between n=2 and n=5. Calculate:

(i) The energy radiation in KJ/mole

(ii) The frequency radiation

(iii)The wavelength of radiation in metres.

Where:

(b) Calculate the ionization energy of hydrogen, given that:

and

(c) The uncertainty in the momentum of a particle is Find its

approximate position

( )

20. (a) The energy of the electron in hydrogen atom in the ground state is given by:

Joules

The energy of the same electron if it occupies a higher level is given by:

Joules

(i) Why is the energy negative?

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(ii) Calculate the energy in joules and the wavelength in metres of the light which

must be absorbed by atom to excite its electron from

(b) The atomic spectrum of hydrogen in the visible region is given by the following

relationship.

(

)

(i) What do symbol, represent

(ii) Calculate the frequency of the third line of visible spectrum.

21. (a) The diagram in the figure below shows the behaviour of the three fundamental particles

when passes through an electric field.

(i) Identify the particles represented by A, B and C.

(ii) Explain the shapes and directions of the path traced by the fundamental particles

as they pass through the electric field.

(b) The emission spectrum of hydrogen consists several series of sharp emission lines in the

ultra visible and infrared regions of the spectrum.

(i) Mention the series which are found in the infrared regions.

(ii) What feature of the electronic energies of the hydrogen atom explains why the

emission spectrum consists of discrete wavelengths rather than a continuum of

wavelength?

(iii)Account for the existence of several series of lines in the spectrum. What quantity

distinguishes one series of liens from another.

(iv) Draw an electronic energy level diagram from the hydrogen atom and indicate on it

the transition corresponding to the line of lowest frequency in the Balmer series.

(v) Explain why the absorption spectrum of atomic hydrogen at room temperature has

only the lines of the Lyman series.

(c) With respect to the electronic configuration justify the following statement:

(i) Electrons are lazy

(ii) Electrons are unfriendly

(iii) Quantum numbers are signature for electrons in the atom.

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22. (a) The diagram below show how a water molecule interacts with hydrogen fluoride

molecule.

(i) What is the value of the bond angle in a single molecule of water

(ii) Explain your answer to part (i) by using the concept of electron pair repulsion

(iii) Name the type of interaction between a water molecules and hydrogen fluoride

molecule shown in the diagram above.

(iv) Explain the origin of the charge shown on the hydrogen atom in the diagram.

(v) When water interacts with hydrogen Fluoride, the value of bond angle in water

change slightly. Predict how the angle is different from that in a single molecule of

water and explain your answer. and hydrazine, are hydrides of elements

which are adjacent in the periodic table.. Data about ethane and hydrazine are

given in the table below:

Melting point /0C -169 +2

Boiling point /0C -104 +114

Solubility in water Insoluble High

Solubility in ethanol High High

(i) Ethene and hydrazine have a similar arrangement of atoms but differently

shape molecules.

What is bond angle in ethane?

What is the Bond angle in hydrazine?

State and explain whether hydrazine is hydrazine and much higher than

those of ethane. Suggest reasons for these differences in term of the

intermolecular forces each compound possesses.

Explain, with the aid of diagram showing Ione pair of electrons and

dispoles, why hydrazine is very soluble in ethanol.

23. (a) Given the name of a geometrical structure and one example of the molecule/compound

formed from the following hybridized atomic orbitals.

(i) hybridised orbital

(ii) hybridised orbital

(iii) hybridized orbital

(b) (i) Give the meaning of Lone pair, Dative bond and sigma bond.

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(ii) By using modern molecular theory complete the table on the following molecules.

Molecule Geometrical structure Name of the structure Hybridisation

(c) The boiling points of water, hydrogen Chloride and argon are shown in the table below:

Substance HCl Ar

Boiling point 0C 100 -85 -186

Total number of electrons 10 18 18

(i) HCl, AR all have Van-der-Waals forces. Outline how Van-der-Waals force

arise beteen molecules

(ii) Liquid has additional intermolecular forces.

What are these forces?

Explain with the aid of diagram, how these forces arise between molecules of

Liquid HCL also has additional intermolecular forces. What are these forces.

Explain the variation in boiling point shown in the above table.

24. (a) Below are elements of group VIA. Their atomic numbers and symbols.

Element Atomic number Symbol

Oxygen 8 0

Sulphur 16 S

Selenium 34 Se

Tellurium 53 Te

(i) Give full electronic distribution in Selenium

(ii) Can Sulphur form S=S? Give reason(s) to support your answer.

(iii) The hydrides of group VIA elements are

(iv) Starting with the least volatile giving clear reasons for your arrangement.

(b) Suggest why the strength of the C-H bond in State the relationship if any, between

the strength covalent bond in and the boiling temperature of and hence state

which one ) has greater boiling point.

(c) State the type of hybridisation shown by the nitrogen atoms in:

(i) (ii) (iii)

25. For each of the following species; identify type of hybridisation of an element in the

bracket, name electron pair geometric shape, molecular geometric shape and bond ange

(i) ( ) (ii) ( ) (iii)

( )

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(a) (i) Predict and explain the polarity of the bonds within (ii) State whether are polar molecules. Explain your answer.

(b) (i) Briefly explain sidgwick – Powel theory

(ii) What type of intermolecular force is found between:

Two water molecules

(c) Although Sodium Potassium are metals, they can never be used in construction of

bridges. Explain.

(d) Starting with the shortest one, arrange the following bonds in order of their bond length.

26. (a) With help of mathematical expression, state the following laws:

(i) Boyles law

(ii) Charles law

(iii) Avogadro’s law

(iv) Dalton law

(v) Graham’s law

(vi) Pressure law

(b) Derive the following equation:

(i) Ideal gas equation

(ii) Equation of a state

(iii)Vander waal equation

(iv) Kinetic theory of a gas equation

(v) Degree of dissociation and association for in terms of density

(c) Describe the characteristics of a gas as influenced by:

(i) Pressure

(ii) Volume

(iii)Temperature

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27. (a) Explain the Kinetic theory of gases and assumption leading to the fundamental kinetic

equation.

(b) Use the fundamental Kinetic equation to deduce the gas laws.

28. (a) Describe molecular masses and relative density of gases.

(b) Describe how the following method are used to find the molecular masses of gases.

(i) Regnaults method

(ii) Dumas method

(iii)Victor Mayer method

(c) Explain the abnormal relative densities of vapour.

29. (a) With examples explain the meaning of colligative properties.

(b) Derive the general equation for all the four (4) colligative propertries.

(c) With mathematical expression state the laws that govern colligative properties.

(i) Raoults law

(ii) Bladgen law

(d) Give practical application of colligative property in daily life.

30. (a) With help of diagram explain the fractional distillation

(b) Give six (6) properties of ideal solution

(c) Differentiate negative deviation from ideal and positive deviation from ideal.

(d) Differentiate the following phrases.

(i) Maximum boiling areotrope and minimum boiling areotrope

(ii) Areotropic mixture and eutedic mixture

(iii)Steam distillation and fractional distillation

(iv) Areotropic point and areotropic mixture

31. (a) Describe the properties of immiscible liquids

(b) Explain the application of immiscible liquid

(c) Using two immiscible liquid A and B show that

(d) Give practical application of steam distillation.

32. (a) State the Nerst distribution law

(b) With help of diagram demonstrate a Nerst distribution law.

(c) Give practical application of distribution law

(d) Explain the deviation from the distribution law

(e) With help of diagram present all the four (4) colligative properties.

33. (a) Explain the difference between osmotic pressure and vapour pressure of a solution.

Calculate the osmotic pressure and vapour pressure of 0.6% aqueous solution of a non-

volatile, non-electrolyte solute, Urea ( ) at The vapour pressure of pure

water at is 24mmHg. Take densities to and assume ideal behavior of the

solution. Gas constant,

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(b) Explain why osmotic pressure method is preferred over depression in freezing point

method for the determination of molar masses of macromolecules.

Calculate the osmotic pressure at and freezing point of 1.8% aqueous solution

( ) Assume ideal behavior of the solution. Take densities to be 1g/mL and for

water to be

(c) Osmotic pressure of a solution containing 7.0g of protein in 100mL of solution is 20mm

Hg at Calculate the molecular mass of the protein.

( )

34. (a) How many mL of a 0.1 M HCl are required to react completely with 1g mixture of

containing equimolar amount of the two?

(b) Calculate the percentage composition in terms of mass of a solution obtained by mixing

300g of a 25% and 400g of a 40% solution by mass.

(c) What role does the molecular interaction play in the solution of alcohol and water?

(d) Why do gases nearby always tend to be less soluble in liquids as the temperature is

raised?

(e) State Henry’s law and mention some of its important applications?

35. (a) Explain why:

(i) Vapour pressures of an aquous solution of sugar is lower than that of pure water.

(ii) Osmotic pressure measurement preferred for determining the molar mass of protein.

(iii)An equimolar solution of sodium Chloride and glucose are not isotonic.

(iv) Glycol-water mixture in a car radiator while driving through the colder region

having sub-zero temperature?

(v) Elevation in boiling point are not used to determine molar mass of protein.

(vi) There is a rise in boiling point when a non volatile solid is dissolved in a liquid.

(b) To of water, of acetic acid is added. If 23% of acetic is

dissociated, what will be the depression in freezing point? Kf and density of water are

respectively?

(c) Two elements A and B form compounds having molecular formulae . When dissolved point by 3.2K, whereas 1g of lowers it by 2.3K. The molar

depression constant for Benzene is Calculate the atomic masses of A

and B.

(d) Calculate the boiling point of a solution containing 0.61g of benzoic acid in 50g of

Assuming 84% dimerization of the acid. The boiling point and Kb of are

and respectively.

36. (a) The molar volume of liquid benzene ( ) increases by a factor of

2750 as it vaporizes at and that of liquid toluene (density ) increases by a factor of A solution of benzene and toluene at has a

vapour pressure of 46.0 torr. Find the molefraction of benzene in the vapour above the

solution.

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(b) 2.0g of benzoic acid dissolved in 25g benzene shows as depression in freezing point

equal to 1.62K. Molar depression constant (Kf) of benzene is What is

the percentage of association of the acid in solution.

37. Define the following terms:

(i) Standard enthalpy of reaction

(ii) Enthalpy

(iii) Standard enthalpy of combustion

(iv) Standard enthalpy of hydrogenation

(v) Standard enthalpy of sublimation

(vi) Dissociation energy

(vii) Atomization energy

(viii) Standard enthalpy of neutralization

(ix) Hess law

(x) Enthalpy change

(xi) Standard state

(xii) Standard enthalpy of formation

(xiii) Bond enthalpy

(xiv) Enthalpy of hydration

(xv) Enthalpy of solution

(xvi) Ionization energy

(xvii) Electron affinity

(b) Draw the Born-haber cycle for:

(i) NaCl

(ii) (iii) (iv)

38. (a) Calculate the standard enthalpy of formation of ethane using the following information:

(i) Given average bond enthalpy:

Atomization enthalpies are

( ) ( )

( )

(b) Calculate the enthalpy change for the reaction:

OH

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Given the mean standard bond enthalpy of the following:

39. (a) Calculate the standard enthalpy of hydrogenation of propene to propane from the data

given below:

Bond Bond Enthalpy ( ) 837

611

347

414

435

(b) Calculate the enthalpy of the formation of methane given that its enthalpy of combustion

is and the enthalpy of formation of water and carbon dioxide are:

respectively.

40. (a) Enthalpies of combustion of some substance and given below: Substance

( )

Hydrogen -242

Benzene -3302

Cyclohexene -3146

Cyclohexane -3940

Calculate enthalpy of hydrogenation of:

(i) Cyclohexene (ii) Benzene (iii) Cyclohexane

(b) Calculate standard enthalpy of combustion of propene from the following information :

41. (a) Using the information below calculate lattice energy of Potassium Bromide.

Atomization energy of K Dissociation energy of

Electron affinity of Ionization energy of

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(b) Carefully study the information given below. Enthalpy changes are given in 1

st ionization energy of ( )

Atomization energy of ( )

2nd

ionization energy of ( )

Heat of atomization energy of oxygen

1st electron affinity of oxygen

2nd

electron affinity of oxygen

Heat of formation of ( )

From the information above, draw a labeled Born Harber cycle for CaO. Calculate the

lattice energy for CaO.

42. (a) A certain reaction between Lithium hydroxide (LiOH) and hydrochloric acid (HCl) take

place in aqueous solution where 50g of solution and causes the temperature of the

solution to increase from The specific heat capacity of water was

. How much heat was supplied and state whether the reason is exothermic

or endothermic?

(b) the mixture was then stirred

rapidly and temperature rose to If density of each solution was

and specific heat capacity was

(i) Explain why solution was mixed rapidly

(ii) Mention any assumption made

(iii)Calculate enthalpy of neutralization

43. (a) Calculate enthalpy of combustion for:

(i) 1g of Carbon (C) ( (ii) 24g of ( (iii)3g of

(b) The equation for the heat of formation of copper (I) oxide ( ) and copper (II) oxide

and respectively. The first and second

ionization energy of copper is 339.3KJ/mole. Draw Born-Haber cycles and calculate the

lattice energy of the two oxides.

44. (a) explain the concept of periodicity

(b) Compare the effort of the following scientist in trying to classify elements

(i) Dobereiner

(ii) Newland octave

(iii)Iother Mayer

(iv) Mendeleev

(v) Moseley

(c) Describe the main features of the modern periodic table.

(d) Compare the periodic law in terms of atomic masses versus atomic number

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45. (a) Describe the general periodic trends in physical properties across periods and down the

group for:

(i) Atomic and ionic size

(ii) Electro negativity

(iii)Ionization energy

(iv) Electron affinity

(b) Describe period tend in physical properties down the group IA, IIA and VII.

(c) Describe periodic trends in physical properties across the periods 3 elements.

46. (a) Describe chemical properties of group I, II and VI with:

(i) (ii) (iii) HCl (iv) (v) and

(b) Describe trend in chemical properties of period 3 elements (hydride, hydroxide,

chlorides and oxide) with water.

47. (a) What is diagonal relationship

(b) State two cause of diagonal relationship.

(c) State four (4) diagonal relationship for the following elements:

(i) Li and Mg

(ii) Be and Al

(iii)B and Si

48. Explain why the first elements in group differ from the rest of a group member using I, II

and VII.

49. (a) Classify metal oxides.

(b) How can you prepare metal oxide?

(c) How metal oxide reacts with:

(i) Water (ii) Dilute acids

(d) Describe reaction of amphoteric oxides.

(e) Explain eight use of metal oxide

50. (a) (i) What is meant by metal hydroxide

(ii) How metal hydroxide is prepared

(b) Describe the properties of metal hydroxide when treated with:

(i) Air (ii) Dilute acid

(c) Explain use of metal hydroxides