THE HALOGENS at a glance

 

HALOGENS

Halogens (salt formers) are found in group VII of the periodic table. They are the most reactive nonmetals. They have seven electrons in their outermost shells and so ionize to form univalent negative ions. They exist as diatomic molecules. They are coloured and they form electrovalent compounds with metals.

 They include are chlorine, fluorine, bromine, iodine and astatine.

ELECTRONIC CONFIGURATION OF HALOGENS 

The halogens have one electron short of the noble gas structure in their  electronic configuration (i.e.  they contains seven electrons in their outermost shells), and the readiness to complete the octet arrangement by receiving an electron makes the halogens very reactive.

The electronic configurations of the halogens are shown below: 

   Fluorine = 9: 1s² 2s² 2p⁵

   Chlorine = 17: 1s² 2s² 2p⁶ 3s² 3p⁵

  Bromine = 35: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁵

 Iodine   = 53: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s² 5p⁵

PHYSICAL PROPERTIES OF THE HALOGENS

1.  They are univalent, and readily accept one electron from other atoms to form ionic compounds (especially from metals e.g Na & K). They also share electrons with themselves or with non-metals to form covalent compounds.

2.   They exist as diatomic molecules.

3.  Fluorine and chlorine are gases, bromine is a liquid and iodine a solid.

4.   The halogens are coloured, with typical penetrating odour. The colours deepen down the group. Fluorine is pale-yellow, chlorine is greenish- yellow, bromine is red and iodine is violet.

5.   They are volatile substances. Their volatility decreases down the group.

6.   All the halogens except fluorine, dissolve to some extent in water, fluorine reacts with water to give oxygen and hydrogen fluoride.

CHEMICAL PROPERTIES OF THE HALOGENS

The halogens are very reactive elements. Their reactivity decreases down the group. Fluorine is the most reactive halogen. They are also strongly electronegative. Their Electronegativity decreases down the group.

1. As oxidizing agents. Halogens are strong oxidizing agent. They readily accept electrons. The oxidizing power decreases down a group.

2.    Reaction with metals: Halogens react with metals to form ionic compounds.

 2Na_(s)  +  F_2(g) →  2NaF_(s)

3.  Reaction with hydrogen: Fluorine explodes with hydrogen even in the dark, chlorine reacts slowly in the dark but explode in bright sunlight, bromine reacts with hydrogen in the presence of platinum catalyst, while iodine reacts partially with hydrogen on heating. Example

              H_2(g) + Cl_2(g) → 2HCl_(g)

Stability of the hydrogen halides decreases down the group. Hydrogen fluoride is a liquid with a boiling point of 19^OC. The other hydrogen halides are gases.        

4.    Reaction with water: Fluorine reacts vigorously with water to give off oxygen gas. 

Chlorine reacts very slowly with water to give a mixture of hydrochloric acid and oxochlorate (I) acid 

             Cl2(g) + H2O → HOCl + HCl

which later decomposes to give hydrochloric acid and oxygen gas.

            HOCl(aq) → HCl(aq) + O2(g)

 The oxygen gas given off by the oxochlorate (I) acid is responsible for the bleaching action of moist chlorine gas and chlorine water.

             H₂O_(g)  +  Cl_2(g)  → HCl_(aq)  +  HOCl_(aq)

CHLORINE

Chlorine is the most important member in the halogen family. It does not occur as free element in nature because it is too reactive. It is usually found in combined state as chlorides.

LABORATORY PREPARATION OF CHLORINE

1.  By the oxidation of concentrated HCl with strong oxidizing agent such as MnO₂ or KMnO₄

             MnO_2(s)+ 4HCl_(aq)→MnCl_2(aq)+ 2H₂O_(l) + Cl_2(g)

                                    

2.         By heating concentrated H₂SO₄ with a mixture of NaCl and MnO₂

                2NaCl_(s) + MnO_2(s) + 2H₂SO_4(aq) →Na₂SO_4(aq) + MnSO_4(aq) + H₂O_(l)+ Cl_2(g)

INDUSTRIAL PREPARATION

Chlorine is manufactured industrially by the electrolysis of brine and molten NaCl, MgCl₂ or CaCl₂

PHYSICAL PROPERTIES

1.   Chlorine is a greenish-yellow gas with unpleasant chocking smell.

2. It is a poisonous gas.

3. It is about 2.5 times denser than air.

4. It is liquefied under a pressure of about 6atm.

5. It is moderately soluble in water.

CHEMICAL PROPERTIES

1.   It is very reactive and forms electrovalent compound with metals and a single covalent bond compounds with non-metals.

2Na_(s)  + Cl_2(g)  →  2NaCl_(s)

Cl_2(g)  +   H_2(g) → 2HCl_(g)

2.  It displaces other halogens from solution of their acids and salts

Cl_2(g)  +   NaI_(aq)  →  2NaCl_(aq)    +    I_2(g)

3. It combines directly with other elements except oxygen, nitrogen carbon and the noble gases; to form chlorides

Ca_(s)   +  Cl_2(g)   → CaCl_2(s)

4. It displaces hydrogen from its compounds due to its strong affinity for hydrogen 

C₁₀H_12(l) + 8Cl_{2(g) →}10C_(s) + 16HCl_(g)

5. It is a powerful oxidizing agent: it oxidizes green Fe²⁺ to yellow or brown Fe³⁺

2FeCl_2(aq) + Cl₂ →2FeCl_3(aq)

6.  It is a bleaching agent:  The bleaching action of chlorine is due to its ability to react with water to form oxochlorate (I) acid which decomposes to release oxygen which in turn oxidizes the dye to form a colourless compound.

H₂O_{(l) +} Cl_2(g)  → HCl_{(aq) +}    HOCl_(aq)

HOCl_(aq)     →   HCl_(aq)  +  [O]

Dye   +   [O] →  [Dye + O]

^{Colored                    Colourless}            

7.  It reacts with hot concentrated NaOH solution to give a mixture of sodiumtrioxochlorate (V) and sodium chloride.

6NaOH  + 3Cl_2(g)    →  NaClO_3(aq)   + 5NaCl_(aq)  +  H₂O_(l)

^{hot concentrated                   Sodium trioxochlorate (V)}

With cold dilute solution of NaOH, a pale yellowish mixture of sodiumoxochlorate (I) and sodium chloride is formed.

2NaOH_(aq)  +  Cl_2(g)  →  NaOCl_(aq)  +  NaCl_(aq)  +  H₂O_(l)

^{cold dilute                             sodium oxochlorate(I)}           

8. It reacts with CaOH solutions to produce bleaching powder

Ca(OH)_2(aq)  +  Cl_2(g) →  CaOCl₂.H₂O_(s)

                                          ^{Bleaching powder}

TEST FOR CHLORINE

1.  It turns wet blue litmus paper pink and then bleaches it.

2.  It turns damped starch-iodide paper blue black.  

Chlorine turns starch-iodide paper blue black because it displaces iodine from the iodide. The iodine liberated then turns the starch blue.

USES OF CHLORINE

1. It is a powerful germicide (due to its oxidizing nature).

2.  It is used as a bleaching agent for cotton, wool, pulp etc.

3.  It is used in the manufacture of polyvinyl chloride (PVC) and synthetic rubber.

4. It is used in the manufacture of organic compounds like  trichloromethan (CHCl_3),  and   tetrachloromethane (CCl₄

5. It is used in producing KClO₃, for making matches and fireworks.

6.  It is used for making NaClO_3, a weed killer.

7.  It is used for making acidified NaClO solution a domestic antiseptic.

COMPOUNDS OF CHLORINE

HYDROGEN CHLORIDE

Hydrogen chloride exists as a gas at room temperature. It dissolves readily in water to form hydrochloric acid. It occurs in traces in the air as industrial by-product and is considered as an air pollutant; but it can be easily washed down as acid rain since it is very soluble in water.

LABORATORY PREPARATION

The gas is prepared by the action of hot concentrated H₂SO₄ on any soluble chloride. Example    2NaCl_(s)  +  H₂SO_4(aq)  →  Na₂SO_4(aq)  +  2HCl_(g)

Note: NaHSO₄ is first formed at a lower temperature and later at higher temperature HCl gas is formed. The gas is dried by passing it through concentrated H₂SO₄ in another flask and collected.

INDUSTRIAL PREPARATION

Pure HCl gas can be produced in large scale by direct combination of hydrogen and chloride gas obtained from the electrolysis of brine.

     H_2(g)+ Cl_2(g)→2HCl_(g)

PHYSICAL PROPERTIES

1.  Pure HCl gas is colourless and has sharp irritating smell

2. It turns damp blue litmus paper red

 3.  It is about 1.25times denser than air

4 It is very soluble in water, hydrochloric acid

5. It dissolved readily in non-polar solvent like chloroform and toluene. 

when HCl is dissolved in nonpolar solvents, the solution does not conduct electricity and has no acidic properties because hydrogen chloride which is a covalent molecule does not ionize when it dissolves in non-polar solvents. But it dissolves in water and ionizes. The ions formed in aqueous solution are responsible for the acidic property and conductivity of its aqueous solution. 

6.  It forms misty fumes in moist air because it dissolves in the moisture to form tiny droplets of HCl acid

CHEMICAL PROPERTIES

1.  Combustion: - Hydrogen does not support combustion 

2.  It combines directly with NH₃   producing dense white fumes of ammonium chloride

 HCl_(g) + NH_3(g) → NH₄Cl_(s)

3. It reacts with electropositive metals to form their respective chloride displacing hydrogen gas 

Zn_(s) +2HCl_(g) → ZnCl_2(s) + H_2(g)

TEST FOR HYDROGEN CHLORIDE

1.  Place a gas rod that has been dipped in ammonia solution over the gas jar containing the unknown gas, if a dense white fumes forms on the glass rod, then the gas is hydrogen chloride gas.

2.   Few drops of silver trioxonitrate (V) is added to the gas jar containing the unknown gas and shaken. If white precipitate of silver chloride is observed, then the gas is hydrogen chloride gas

CHLORIDES

Chlorides are normal salts formed when metallic ion replace the hydrogen ion in hydrochloric acid. Soluble Chlorides are prepared by neutralization reaction while insoluble are prepared by double decomposition method.

All Chlorides are soluble in water with except, AgCl, HgCl2, PbCl2

PROPERTIES

1.  Chlorides are stable to heat, that is, they are not decomposed by heat.  

They are recovered from solution by evaporation to dryness and sometimes by crystallization.

2. They react with hot concentrated tetraoxosulphate (VI) acid to produce hydrogen chloride gas.

2NaCl_(s)+H₂SO_4(aq) →Na₂SO_4(aq)+2HCl_(g)

and in the presence of a strong oxidizing agent, chlorine is produced.

ZnCl_2(s) + KMnO_4(s)+ 2H₂SO_4(aq) → ZnSO_4(aq)+ K₂SO_4(aq) + 2MnO_2(aq) + 2H₂O_(l)+Cl_2(g)

TEST FOR CHLORIDES 

Add some Few drops of AgNO_3(aq) to a solution of the sample in a test tube, if it forms   a white precipitate, now acidify the solution by adding dilute trioxonitrate acid if the white precipitate remains but readily dissolves in excess NH_3(aq) solution then a chloride ion is present.

OBJECTIVE QUESTIONS

1. The bleaching action of chlorine is through the process 

a. Hydration 

b. Hydrolysis 

c. Reduction 

d. Oxidation 

2. When chlorine is passed through a sample of water, the pH of the water sample would be 

a. <7

b. =7

c. >7

d. 0

3. Halogens generally react with metals to form 

a. Alkalis.

b. Acids.

c. Bases.

d. Salts.

4. Potassium chloride solid does not conduct electricity because 

a. It is a covalent compound 

b. Stong cohesive forces make its ions immobile 

c. Strong cohesive forces make its molecules immobile 

d. Each of Potassium and chlorine ions has a noble gas structure.

5. Chlorine water is used as a bleaching agent because it is 

a. An acidic solution 

b. An alkaline solution 

c. An oxidizing agent 

d. A reducing agent 

6. Which of the following halogens is a liquid at room temperature 

a. Iodine

b. Chlorine 

c. Bromine

d. Fluorine.

7. Chlorine, bromine and iodine belong to the same group and 

a. Are gaseous at room temperature 

b. Form whit precipitate with AgNO3(aq)

c. React violently with hydrogen without heating.

d. React with alkali 

8. Which of the following chlorides is insoluble in water? 

a. AgCl

b. KCl

c. NH4Cl

d. ZnCl2

9. Which of the following statements about chlorine and iodine at room temperature is correct 

a. Chlorine is a gas and Iodine as solid 

b. Chlorine is a liquid and iodine is a gas.

c. Chlorine and iodine are gases 

d. Chlorine is solid and iodine is a liquid.

10. 

THEORY QUESTIONS

1. Draw and label a diagram to illustrate the preparation and collection of a dry sample of chlorine gas in the laboratory. 

1b. state two use of chlorine 

2a. Write the equation for reaction between chlorine gas and            

  i. Concentrated NaOH

   Ii. Dilut NaOH

b. Write the electronic configuration of the following atoms/ions: Cl, F^-, Br.

C.  Give three physical properties of the halogens

2a.   Explain one laboratory preparation of dry chlorine gas.

b. Name the method of collection of chlorine gas and explain why it can be collected by the method. 

3a.   Mention three physical properties of chlorine.  

b.  Using balanced equations, state THREE chemical properties 

C. Explain why hydrogen chloride in toluene does not conduct electricity but its aqueous solution  conducts  electricity.

4a.   Describe a test for a soluble chloride.

bGive three uses of chlorine gas.

C.  State TWO physical and TWO chemical properties of hydrogen chloride gas

5a.  An unknown gas is colourless, has an irritating smell, fumes in moist air and turns blue litmus paper red; describe how you will confirm the gas to be hydrogen chloride gas.

b. A solid chloride E which sublimed on heating reacted with an alkali F to give a choking gas G. G turned moist red litmus paper blue.  Identify E,F and G. 

PERIODIC TABLE at a glance

PERIODIC TABLE

 The periodic table is an arrangement of all the elements in a particular order.

Mendeleev’s Periodic Table (1869): Dmitri Mendeleev created the first widely accepted periodic table. He arranged elements by atomic mass and left gaps for undiscovered elements, predicting their properties accurately.

Modern Periodic Table (1913): Henry Moseley improved the table by arranging elements according to atomic number instead of atomic mass, resolving earlier inconsistencies.

Modern Periodic Law

The Modern Periodic Law states that:

“The physical and chemical properties of elements are a periodic function of their atomic numbers.”

This means:

i. When elements are arranged in order of increasing atomic number, their properties repeat at regular intervals.

ii. Elements in the same group (vertical column) have similar chemical properties because they have the same number of valence electrons.

        Arrangement of the first 20 elements 

  I           II                  III       IV    V    VI      VII    VIII

 

1H								2He
3Li	4Be		5B	6C	7N	8O	9F	10Ne
11Na	12Mg		13Al	14Si	15P	16S	17Cl	18Ar
19K	20Ca

     → PERIOD  

            ↓

         GROUP

Each horizontal row is called a period (I - VIII)

while the vertical column is called a group

The periodic table and the electronic configuration: -The largest principal quantum number of the electronic configuration of an element (the highest positive integer) represents the period to which an element belongs to while the number of electrons in the outermost shell of the atom represents the group to which the element belongs to in the periodic table. For example, 

Given two elements X and Y with the following electronic configuration 

i. X=1s²2s²2p⁴ and element  

ii.  Y = 1s²2s²2p⁶3s². 

PERIOD

The largest number (positive integer = principal quantum number) in X is 2 (black bold) (i.e X contains 2 shells) and hence belongs to period 2. The largest number (principal quantum number) in Y is 3 (i.e Y contains 3 shells) and belongs to period 3. Simply put the number of electronic shells in an atom is equivalent to its period in in the Periodic Table.

GROUP

The total number of electrons in the outermost shell of X is 6 i.e  (2+4) and so it belongs to group 6 in the periodic table while Y belongs to group 2 (as it has only 2 electrons in its outermost shell).

TRENDS/ PERIODICITY IN THE PERIODIC TABLE

Periodicity is the variation of properties of elements as you move across a period from left to right or as you go down a group.

These properties are also known a trends in the periodic table and they vary in intensity as you move across the period from group 1 to group 8 and down the group from top to bottom

These properties include: -

ATOMIC RADIUS: - This is the size of an atom. It is the distance between the nucleus of atom and the outermost shell.

It decreases across the period and increases down the group in the periodic table. 

Reason

Across the period as the atomic number increases the charge on the nucleus (nuclear charge) also increases, since the electrons are entering into the same shell, they will experience a greater attraction pulling them towards the center of the atom and hence a decrease in size of the atom across the period. But down the group new shells are being added and hence the atomic size increases automatically.

 

.

        Size of the atoms decreases as you move across the period but increases down the group

IONIC RADIUS: -For metals their atomic radius is larger than their ionic radius this is because metals ionize by the loss of the outermost or valence electrons and so the ion becomes one shell less than the atom.  Hence the smaller ionic radius.

                                 Atomic Radius vs Ionic Radius

                   
       here the sodium atom is larger in size than the sodium ion due to the loss of the outermost electron/shell. Similarly, the atomic radius of magnesium is smaller than the ionic radius of the magnesium ion

For non-metals their atomic radius is smaller than their ion radius, since non-metals ionize by gaining electros. A slight repulsion occurs between the gained electron and the other electrons in the valence shell. This results to a slight expansion of the ionic radius.     

              

    here the chlorine atom is smaller than the chloride ion due to the repulsion between the valence electrons and the gained electrons. Similarly, the atomic radius of sulphur atom is smaller the ionic radius of the sulphide ion

IONIZATION ENERGY: - This is the energy required to remove a valence electron from an atom in the gaseous state to form a mole of gaseous ions.

 It increases across the period (due to an increase in the nuclear attraction on the valence electrons across the period) and decrease down the group (as the valence electrons get farther away from the nucleus the become less attracted to the nucleus)

           

         Ionization Energy of the elements on the Periodic Table

                                

ELECTRONAGATIVITY: - This is the tendency of an atom to attract electrons to itself in a molecule. It increases across the period and decrease down the group  

                                 

                          The electronegativities of the elements in the Periodic Table

ELCTRON AFFINITY: - This is the energy liberated when an electron enters an atom in the gaseous state to form a mole of negative ion. It increases across the period and decreases down the group.

ELECTRICAL CONDUCTIVITY: - Sodium, magnesium and aluminum are good conductors of electricity because of the ‘sea’ of delocalized electrons they possess. Silicon is a semi-conductor, but not as good a conductor as graphite. All the other elements are electrical insulators.

GENERAL PROPERTIES OF ELEMENTS IN EACH GROUP

1.     GROUP I (s-block elements) (Alkali metals) 

             (Li, Na, K, Rb, Cs and Fr)

i.    They are soft, malleable, and ductile

ii.   They ionize by loss of one electron

iii.  They are good reducing agents

iv.   They are good conductors of heat and electricity

v.    Their densities  generally increase down the group  

vi.    They react with cold water to displace hydrogen gas

            Na(s) + H2O(l) → NaOH(aq) + H2(g)

2.     GROUP II (s-block) (Alkaline earth metals)

         (Be, Mg, Ca, Sr, Ba and Ra)

i.    They ionize by the loss of two electrons 

ii.   They are good conductors of heat and electricity

iii.   They are good reducing agents ( because they lose electrons readily)

iv.    Their melting and boiling points decreases generally down the group

v.      Their densities increases down the group

`3.     GROUP III (p-block)( The boron family)

         (B, Al, Ga, In and Ti)

  i.     Apart from boron all other members of the group are metals 

  ii.   They ionize by losing 3-electrons (common oxidation state is +3)

  iii.   Boiling point decreases down the group (but increases across the period

  iv.   They have high melting points

  vi.   They all form oxides when strongly heated in oxygen

  vii.     Thier reactivity increase down the group

 viii. They tarnish readily in air due to the formation of an oxide layer

4.       GROUP IV (p-block elements) (The Carbon family)

           (C, Si, Ge, Sn and Pb)

 i.      They have oxidation states of +2 and +4 but the +2 becomes more common

 ii.     C (non-metal) Si and Ge (metalloid) have covalent /bonding network within the network while                 Sn and Pb are metallic

 iii.     Their oxides range from acidic (CO2) to amphoteric (SiO2)  

 Iv    

5.      GROUP V (p-block elements) (The Nitrogen family)

                (N, P, As, Sb and Bi)

 i.  Members exhibit various oxidation state but as you go down the group the +3 

oxidation state becomes predominant

 ii.     There is a gradual change in the properties of the members of the group moving from individual or single molecules (N and P) to covalent networks (As and Sb) to metal (Bi)

6.       GROUP VI (p-block)( The Oxygen family)

 The elements in this group and their electronic configuration are shown below

Oxygen = 8: - 1s² 2s² 2p⁴

Sulphur = 16: - 1s² 2s² 2p⁶ 3s² 3p⁴

Selenium= 34: - 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁴

Tellurium = 52: - 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s² 5p⁴

Polonium = 84: - 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s² 5p⁶ 5d¹⁰ 5f¹⁴ 6s² 6p⁴

GENERAL PROPERTIES OF GROUP VI ELEMENTS

The elements in this group include Oxygen, Sulphur, Selenium, Tellurium and Polonium.  

1.   Every other element except oxygen is a solid at room temperature 

2.  Metallic property increases down the group. (Oxygen and sulphur are non-metal; selenium and tellurium are metalloid, while polonium is a metal).

3. Oxygen and sulphur exhibit allotropy.

4.   With six electrons in their outermost shell. Hence their oxidation number is -2; though sulphur can displays -4 and -6 states in some compounds.

5.  Electronegativity decreases down the group. Thus, oxygen is a good oxidizing agent.

6. They are good oxidizing agent (because they readily accept electrons)

7.  They do not conductor electricity

8.  They do not react with water

 

7.      GROUP VII: - (p-block) (Halogens)

i.  They ionize by gaining one electron

ii.  They are good oxidizing agents

iii.  They are coloured

         * Florine is yellowish 

         * Chlorine is greenish yellow 

         * Bromine is reddish-brown 

         * Iodine is violet

iv.  They dissolve in water to produce acids 

8.        GROUP VIII or 0 (Noble gases) (rare gases) (inert gases)

             (He, Ne, Ar , Kr, Xe, Rn)

i.  They exist freely as monoatomic molecules in the atmosphere, 

ii. They have no bonding electrons in the outermost shell.

iii. They are non-reactive elements, because their valence shell is completely filled.

iv. They exhibit similar properties among themselves.

v. They bear no resemblance to the halogens that come before them and the alkali metals that come after them. 

vi. Their melting and boiling points increase down the group

 vii. Their ionization energy decreases down the group from helium to radon.

TRANSITION METALS: (d-block elements)

     Transition metals are metals that have partially filled d-orbitals. These elements lie between group 2 and 3 from period 4 in the periodic table. They are metals with special properties. 

  Characteristics of transition elements

i. They have variable oxidation states

ii. They form complex ions

iii.  They form coloured ions

iv.  They are paramagnetic

 v. They are mainly used as catalysts

LANTHANIDES AND THE ACTINIDES

OBJECTIVE QUESTIONS 

  Use the following portion of the periodic table to answer questions 1 to 3

1. Which of the letters indicate elements which exist as diatomic gases.

a).  B and G

b).  A and  F

c).  C and A

d).  A and E

2. Which of the letters represents an alkaline earth metal?

a).  F

b).  E

c).  D

d).  C

3. Which of the following pairs of letters denotes elements containing the same number of electrons in their outermost shells?

a).  C and D

b).  E and F

c).  B and G

d). A and B

4. An element X has electronic configuration 1s²2s²2p⁶3s²3p⁶4s². To which group of the periodic table does X belong?

(a). I   (b). II           (c). III           (d). IV

 

5. Which of the following sets of elements is arranged in order of increasing first ionization energy?

a). ₁₁Na, ₃Li, ₁₉K, ₃₇Rb

b). ₃₇Rb, ₁₉K, ₃Li, ₁₁Na

c).  ₃Li, ₁₉K, ₁₁Na, ₃₇Rb

d). ₃₇Rb, ₁₉K, ₁₁Na, ₃Li

6. Elements which belongs to the same group in the periodic table are characterized by

a). difference of +1 in the oxidation numbers of successive members 

b). Presence of the same number of outermost electrons I the respective atoms

c). difference of 14 atomic mass units between successive members 

d). presence of the same number of electron shells in the respective atoms.

7. Which of the following electronic configuration represents that of a noble gas 

a). 2,8,8,2

b). 2,8,2

c). 2,8

d). 2,6

8. Which of the following pairs of species contains the same number of electrons [ ₆C, ₈O, ₁₀Ne, ₁₁Na, ₁₂Mg ₁₃Al, ₁₇Cl]

a). Mg²⁺ and Al³⁺

b). Cl^- and Ne

c). Na⁺ and Mg

d). C and Cl^-

9. Which of the following statements about rare gases are correct? 

I. Their outermost shells are fully filled.    II. They are generally unreactive.    III. Their outermost shells are partially filled.    IV. They lone pairs of electrons in their outermost shell.

a).  I and II only 

b). II and III only

c). I, II and III only

d). I, II, III and IV

10. How many electrons are in the ion F^- ? [¹⁹₉F]

a). 8      (b) 9      (c). 10      (d) 19

11. Which of the following of properties of elements generally increase down a group in the periodic table?

a). Electron affinity 

b). Electronegativity

c). Ionic radius 

d). Ionization energy

12. In which of the following atoms is the ionic radius larger than the atomic radius?        [₁₁Na, ₁₂Mg, ₁₃Al, ₁₇Cl]

a). Aluminum

b). Chlorine

c). Magnesium

d). Sodium

13. Which of the following properties is characteristics of the halogens?

a). Ability to accept electrons readily.

b). Ability to donate electrons readily.

c). Ability to form basic oxides. 

d). Formation of coloured compounds.

14. 

THEORY QUESTIONS 

1. The electronic configuration of five elements represented by the letters P, Q, R, S and T are indicated below

P --- 1s²2s²2p²

Q --- 1s²2s²2p⁴

R --- 1s²2s²p⁶

S --- 1s²2s²2p⁶3s²

T --- 1s²2s²2p⁶3s²3p⁵

 Without identifying the elements, state which of them

i).  Belongs to group VI in the periodic table

ii).  Is strongly metallic in character

iii).  Readily ionizes by gaining one electron

iv).  Contains two unpaired electrons in the ground state atom.

v).   Readily loses two electrons during chemical bonding

vi).  Does not participate readily in chemical reactions

vii).   Is an s-block element

bi). Copy and complete the table below as appropriate

Particle	Number of Protons	Number of Electrons	Number of Neutrons
11H	1	1
2713Al3+
168O			8

ii). Give the reason why atomic radius increases down a group in the periodic table but decreases from left to right.

iii). State three properties of transition element. [waec]

2. The electronic configuration of atoms of elements A, B, C and D are given as follows

a). 1s²2s²2p²

b). 1s²2s¹

c).  1s²2s²2p⁶

d). 1s22s22p1

ai.  Arrange the elements in order of increasing atomic size, giving reason

ii).  State which of the elements

  I. is divalent 

II. Contains atom with two unpaired electrons in the ground state.

III). Readily loses one electron from its atom during chemical bonding

IV)  Belongs to group III in the Periodic Table.

2(a)(i). List three properties of elements which increases generally across a period in the periodic table. 

(ii). Explain briefly why there is general increase on the first ionization energies of the elements across the period in the periodic table 

HYDROCARBONS at a glance

HYDROCARBONS AND BASIC TERMINOLOGIES IN ORGANIC CHEMISTRY

Introduction

Organic chemistry is the branch of chemistry that deals with carbon-containing compounds, particularly those containing carbon-hydrogen (C-H) bonds. The vast majority of organic compounds are derived from hydrocarbons, making them the foundation of organic chemistry.

Hydrocarbons are important sources of fuel, raw materials for industries, and building blocks for numerous synthetic products such as plastics, pharmaceuticals, detergents, dyes, and fertilizers.

Basic Terminologies in Organic Chemistry

Before studying hydrocarbons, it is important to understand some common terms used in organic chemistry.

1. Organic Compounds

Organic compounds are compounds that contain carbon atoms bonded to hydrogen atoms and may also contain oxygen, nitrogen, sulfur, phosphorus, or halogens.

Examples:

i. Methane (CH₄)

ii. Ethanol (C₂H₅OH)

iii. Acetic acid (CH₃COOH)

2. Hydrocarbons

Hydrocarbons are organic compounds made up of only two elements, carbon and hydrogen.

Examples:

i. Methane (CH₄)

ii. Ethane (C₂H₆)

iii. Benzene (C₆H₆)

Hydrocarbons are classified into:

1. Aliphatic hydrocarbons

2. Aromatic hydrocarbons

3. Homologous Series

A homologous series is a family of organic compounds having the same functional group and general formula, with successive members differing by a CH₂ group

Examples:

Compound	Molecular Formula
Methane	CH₄
Ethane	C₂H₆
Propane	C₃H₈
Butane	C₄H₁₀

Characteristics:

i. Similar chemical properties

ii. Gradual change in physical properties

iii. Same functional group

iv. Successive members differs by a CH₂

4. Functional Group

A functional group is an atom or group of atoms responsible for the characteristic chemical reactions of an organic compound.

Examples:

Functional Group	Name
–OH	Alcohol
–COOH	Carboxylic Acid
–CHO	Aldehyde
–NH₂	Amine

5. Isomerism

Isomerism is the phenomenon where compounds have the same molecular formula but different structural arrangements.

Example:

C₄H₁₀ exists as:

i. n-Butane

ii. Isobutane

6. Saturated and Unsaturated Compounds

Saturated Compounds

Contain only single carbon-carbon bonds.

Example:

i.Ethane (C₂H₆)

Unsaturated Compounds

Contain at least one double or triple bond.

Examples:

i. Ethene (C₂H₄)

ii. Ethyne (C₂H₂)

7. Molecular Formula

Shows the actual number of atoms present in a molecule.

Example:

i. Ethane = C₂H₆

8. Structural Formula

Shows how atoms are connected within a molecule.

Example:

CH₃–CH₃

Hydrocarbons

Hydrocarbons are the simplest organic compounds and form the basis of all organic chemistry.

Sources of Hydrocarbons

1. Crude oil (Petroleum)

2. Natural gas

3. Coal

4. Biomass

Classification of Hydrocarbons

Hydrocarbons are broadly classified into:

1. Aliphatic Hydrocarbons

These consist of straight-chain, branched-chain, or non-aromatic cyclic compounds.

They are divided into:

A. Alkanes

Alkanes are saturated hydrocarbons containing only single covalent bonds.

General Formula

CₙH₂ₙ₊₂

Examples

Alkane	Formula
Methane	CH₄
Ethane	C₂H₆
Propane	C₃H₈
Butane	C₄H₁₀

Properties

• Relatively unreactive

• Undergo combustion

• Undergo substitution reactions

Uses

• Domestic cooking gas

• Fuel for vehicles

• Industrial heating

B. Alkenes

Alkenes are unsaturated hydrocarbons containing at least one carbon-carbon double bond.

General Formula

CₙH₂ₙ

Examples

Alkene	Formula
Ethene	C₂H₄
Propene	C₃H₆
Butene	C₄H₈

Properties

• More reactive than alkanes

• Undergo addition reactions

Uses

• Manufacture of plastics

• Production of alcohols

• Chemical synthesis

C. Alkynes

Alkynes are unsaturated hydrocarbons containing at least one carbon-carbon triple bond.

General Formula

CₙH₂ₙ₋₂

Examples

Alkyne	Formula
Ethyne	C₂H₂
Propyne	C₃H₄

Properties

• Highly reactive

• Undergo addition reactions

Uses

• Welding and cutting metals

• Production of industrial chemicals

Aromatic Hydrocarbons

Aromatic hydrocarbons contain one or more benzene rings.

Examples

• Benzene (C₆H₆)

• Toluene (C₇H₈)

• Naphthalene (C₁₀H₈)

Characteristics

• Highly stable

• Possess delocalized electrons

• Undergo substitution reactions

Uses

• Solvents

• Dye manufacture

• Pharmaceutical production

Reactions of Hydrocarbons

1. Combustion

Hydrocarbons burn in oxygen to produce carbon dioxide and water.

Example:

CH₄ + 2O₂ → CO₂ + 2H₂O

Uses:

i. Energy generation

ii. Domestic cooking

iii. Transportation

2. Substitution Reaction

Common in alkanes and aromatic compounds.

Example:

CH₄ + Cl₂ → CH₃Cl + HCl

3. Addition Reaction

Occurs in alkenes and alkynes.

Example:

C₂H₄ + H₂ → C₂H₆

4. Polymerization

Small molecules combine to form large molecules called polymers.

Example:

n(C₂H₄) → (–CH₂–CH₂–)ₙ

Polymer produced: Polyethene

Importance of Hydrocarbons

1. Source of energy and fuel.

2. Raw materials for petrochemical industries.

3. Manufacture of plastics and synthetic fibres.

4. Production of detergents and solvents.

5. Pharmaceutical manufacturing.

6. Production of fertilizers and pesticides.

Environmental Effects of Hydrocarbon Use

Advantages

i. High energy content

ii. Easily available

iii. Versatile industrial applications

Disadvantages

i. Air pollution

ii. Greenhouse gas emissions

iii. Global warming

iv. Oil spill contamination

Conclusion

Hydrocarbons are the fundamental compounds of organic chemistry and serve as the building blocks for countless organic substances. Understanding their classification, properties, reactions, and applications provides a strong foundation for studying advanced organic chemistry. Knowledge of key organic chemistry terminologies such as functional groups, homologous series, isomerism, and saturation is essential for understanding the behavior and reactions of organic compounds.

Objectives Questions 

Instructions: Choose the correct option (A–D) for each question.

1. Organic chemistry is the branch of chemistry that deals mainly with the compounds of: A. Oxygen
B. Carbon
C. Nitrogen
D. Sulphur

2. The simplest hydrocarbon is: A. Ethane
B. Methane
C. Propane
D. Butane

3. Hydrocarbons are compounds containing only: A. Carbon and oxygen
B. Carbon and nitrogen
C. Carbon and hydrogen
D. Hydrogen and oxygen

4. Which of the following belongs to the homologous series of alkanes? A. C₂H₄
B. C₃H₆
C. C₄H₁₀
D. C₂H₂

5. The general formula of alkanes is: A. CₙH₂ₙ
B. CₙH₂ₙ₋₂
C. CₙH₂ₙ₊₂
D. CₙHₙ

6. Which of the following is an alkene? A. CH₄
B. C₂H₆
C. C₂H₄
D. C₃H₈

7. The functional group present in alkenes is: A. Triple bond
B. Single bond
C. Double bond
D. Hydroxyl group

8. The general formula of alkenes is: A. CₙH₂ₙ
B. CₙH₂ₙ₊₂
C. CₙH₂ₙ₋₂
D. CₙHₙ

9. Which of the following is an alkyne? A. Ethene
B. Ethane
C. Ethyne
D. Propene

10. The general formula of alkynes is: A. CₙH₂ₙ
B. CₙH₂ₙ₊₂
C. CₙH₂ₙ₋₂
D. CₙHₙ

11. The IUPAC name of CH₄ is: A. Methane
B. Ethane
C. Propane
D. Butane

12. The IUPAC name of C₂H₆ is: A. Methane
B. Ethane
C. Ethene
D. Ethyne

13. The first member of the alkene series is: A. Methene
B. Ethene
C. Propene
D. Butene

14. The process by which large hydrocarbon molecules are broken into smaller ones is called: A. Polymerization
B. Fractional distillation
C. Cracking
D. Hydrogenation

15. The main constituent of natural gas is: A. Ethane
B. Propane
C. Butane
D. Methane

16. The test used to distinguish an alkene from an alkane is: A. Litmus test
B. Bromine water test
C. Flame test
D. pH test

17. An alkene decolourizes bromine water because it: A. Is acidic
B. Contains a double bond
C. Is basic
D. Contains oxygen

18. Which of the following hydrocarbons has the highest carbon content? A. CH₄
B. C₂H₆
C. C₃H₈
D. C₄H₁₀

19. The major products of complete combustion of hydrocarbons are: A. CO and H₂O
B. CO₂ and H₂O
C. C and H₂O
D. CO₂ and H₂

20. The brown colour of bromine water disappears when shaken with: A. Methane
B. Ethane
C. Ethene
D. Propane

21. Which of the following is an aromatic hydrocarbon? A. Ethane
B. Propene
C. Benzene
D. Ethyne

22. The molecular formula of benzene is: A. C₆H₁₂
B. C₆H₆
C. C₆H₁₄
D. C₆H₁₀

23. The functional group present in alcohols is: A. –COOH
B. –OH
C. –CHO
D. –NH₂

24. The IUPAC name of CH₃CH₂OH is: A. Methanol
B. Ethanol
C. Propanol
D. Butanol

25. The homologous series characterized by the –COOH group is: A. Alcohols
B. Alkanes
C. Carboxylic acids
D. Esters

26. Which of the following compounds is a carboxylic acid? A. Ethanol
B. Ethanoic acid
C. Ethene
D. Ethanal

27. The process of joining many alkene molecules together to form a giant molecule is called: A. Cracking
B. Hydrogenation
C. Polymerization
D. Distillation

28. PVC and polythene are examples of: A. Fuels
B. Polymers
C. Alcohols
D. Acids

29. The source of most hydrocarbons used as fuels is: A. Limestone
B. Air
C. Crude oil
D. Water

30. Which of the following compounds is unsaturated? A. Methane
B. Ethane
C. Propane
D. Ethene

Rates of Chemical Reactions

During a chemical reaction, reactants collide with one another to form products, and the formation of these products do not occur at the same rates. Hence 

Rate of a chemical reaction can be defined as the number of moles of reactants that are converted, or products that are formed per unit time.

Mathematically 

Rate =      mass in grammes 

                      time taken

For instance, if 6g of zinc metal is placed in dilute tetraoxosulphate (VI) acid and it takes 3 mins to completely react, then the rate of chemical reaction is given by. 

Rate =      6g     = 2g/mins

                3min

That is, 2g of the zinc was converted to ZnCl₂ per minute 

Factors used for Measuring Rates of Reaction

The rate of a chemical reaction can be measured or can be determine by any one of the following.

i.   decrease in mass of reactants.

ii.  increase in volume of a gaseous product 

iii. change in pH

iv. change in colour intensity 

v. Change in pressure

 

These are all measurable factors.

 Factors Affecting the Rate of Chemical Reactions.

The following factors will a reaction to be fast or slow.

i. Nature of reactants

ii. Temperature

iii.. Concentration 

iv. Pressure 

v.  Presence of light

vi. Surface area 

vii. Presence of a catalyst

Before we discussed how each of these factors will affect the rate of a chemical reaction. It is important that we understand the concept of collision theory.

Collision theory assumes that for a chemical reaction to occur, there must be Collision between reactants particles, and these Collisions must be effective.

A collision is said to be effective when it leads to formation of products

What this concept is actually implying is that all collisions do not lead to the formation of a product, only the ones that are effective. So, the various factors that affect chemical reactions are factors that actually increase the number of effective collisions.

I. Effect of Nature of reactants: - this is one of the factors that affects the rate of a reaction, since different elements/ compounds behave differently.

The reactivity of metals varies as you go down the activity series and so when metals react with acids for example the reactions are not the same.

Example by virtue of their nature (their high reactivity) sodium and potassium will react explosively with dilut acids while metals like zinc and iron will react moderately with dilute acids.

II. Effect of Temperature: - An increase in temperature will cause reactants particles to gain more kinetic energy leading to an increase in effective collision and hence, an increase in the rate of reaction. An increase in temperature will also lead to an increase in the energy of the system.

III. Effect of Concentration: - An increase in the concentration of the reactants will lead to an increase in the number of reactants per unit area (the reactants becomes closer) thus, leading to overcrowding and increase in the effective collision of the reactants and hence an increase in the rate of the reaction.

IV. Effect of pressure: - An increase in pressure will lead to a decrease in volume that is, a decrease in the intermolecular space between the reactants particles leading to an increase in the effective collision and hence an increase in the rate of reaction. And vice versa 

V. Effect of presence of light: - Some reactions are photochemical, that is, affected by light. When such reactions are exposed to light the reactants, particles become more activated and collide more increasing the number of effective collisions and hence an increase in the rate of reaction.

VI. Effect of a catalyst: - A catalyst increases the rate of a reaction by creating a different pathway with a lower activation energy.

ACTIVATION ENERGY: see definition in chemistry.

  For a collision to be effective, the reactants must possess the minimum amount of energy needed to overcome the energy barrier (activation energy for the reaction) for that reaction. The higher the activation energy, the slower the rate of reaction; positive catalyst helps to lower the activation energy.

 Energy Profile Diagram 

Energy profile diagram is a diagram / graph that shows the pathway of a chemical reaction. whether the reaction is exothermic or endothermic.

                                          Figure1: Energy profile diagram of an exothermic reaction. 

The reactants have higher energy than the products, showing that heat is released during the reaction. The peak represents the activated complex, while Ea is the activation energy and ΔH is the heat of reaction.

                                         Figure2: Energy profile diagram of an endothermic reaction.

 The products have higher energy than the reactants, indicating that heat is absorbed during the reaction. The highest point is the activated complex, Ea is the activation energy, and ΔH represents the heat of reaction.

 Rate Curve

A rate Curve is the graph which shows the rate of a reaction. It is a graph of reaction against time.( that is, change in concentration against time, decrease in mass against time, e.t.c) 

The slope or the gradient of the curve is steep at the beginning because the reaction is fastest ( since  , it becomes less steep as the reaction progresses and slows down, then it finally becomes horizontal. 

The point at which the graph becomes horizontal indicates the end point of the reaction, when one of the reactants is completely used up

figure 1:  A typical rate curve showing how the mass of a reactant decreases with time during a chemical reaction. The steep part of the curve shows a fast reaction at the beginning, while the flatter part shows the reaction slowing down as the reactants are used up.

Figure 2: A graph showing how the volume of a gaseous product increases with time during a chemical reaction. The curve rises steeply at first, indicating a fast rate of gas production, and then levels off as the reaction nears completion.

ORDER Of REACTIONS

1. A first-order reaction is a chemical reaction in which the rate depends on the concentration of only one reactant raised to the power one.

In simple terms:
The speed of the reaction is directly proportional to the concentration of one reactant.

           Rate ɑ [A]
   or

           {Rate} = k[A]

Where:

• (A) = reactant

• (k) = rate constant

What this means

If the concentration of the reactant is doubled, the rate of reaction also doubles.
If the concentration is halved, the rate is halved.

Example

The decomposition of hydrogen peroxide:

H₂O₂ → H₂O + O₂

The rate depends only on the concentration of (H₂O₂), so it is a first-order reaction

A first-order reaction has a constant half-life — the time taken for half of the reactant to be used up does not change, no matter the starting concentration.

Example of a First-order reaction)

The half-life of a first-order reaction is 10 minutes.
If the initial concentration of the reactant is 80 mol dm⁻³, find the concentration after 20 minutes.

Solution

For a first-order reaction, the half-life is constant.

In 10 minutes → concentration becomes half

                 80 →40

In another 10 minutes (total = 20 minutes) → it halves again

            40 → 20

Answer

The concentration after 20 minutes is 20 mol dm⁻³.

A second-order reaction is a chemical reaction in which the rate depends on the square of the concentration of one reactant, or on the product of two reactants.

It can be written as:

                                         Rate = k[A]²
                             or 
                          

                         Rate = k[A][B]

what this means is that

• If the concentration of a reactant is doubled, the rate increases four times.

• If the concentration is tripled, the rate increases nine times.

Simple example

The reaction between nitrogen dioxide molecules:

                           2NO₂ →2NO + O₂

This is second order because the rate depends on ([NO₂]²).

Another example of a 2nd order reaction is 

The reaction between hydrogen and iodine:

                                H₂ + I₂ →2HI

This reaction is second order because experiments show that its rate depends on the product of the concentrations of both reactants:

Rate = k[H₂][I₂]

It means:

• If the concentration of H₂ is doubled, the rate doubles.

• If the concentration of I₂ is also doubled, the rate doubles again. So the overall rate becomes four times faster.

The reaction is second order because the rate depends on the concentrations of two reacting substances.

waec/neco keypoint

A second-order reaction does not have a constant half-life — its half-life changes as the concentration changes.

OBJECTIVE QUESTIONS 

1. Which of the following best describes the rate of a chemical reaction?

A. The amount of heat produced
B. The speed at which reactants are used up
C. The colour change in a reaction
D. The total energy of the reaction

2. Which of the following factors does NOT affect the rate of a chemical reaction?

A. Temperature
B. Pressure
C. Concentration
D. Density

3. An increase in temperature increases the rate of reaction because

A. molecules become heavier
B. molecules move faster and collide more
C. the volume of the gas increases
D. the pressure decreases

4. Which of the following will increase the rate of reaction between zinc and hydrochloric acid?

A. Using dilute acid
B. Using powdered zinc
C. Lowering the temperature
D. Using large zinc pieces

5. A catalyst increases the rate of a chemical reaction by

A. increasing the temperature
B. creating a different pathway with a lower activation energy
C. increasing the pressure
D. increasing the concentration

6. Which of the following is an example of a catalyst?

A. Copper (II) sulphate
B. Manganese (IV) oxide
C. Sodium chloride
D. Water

7. In a reaction between a solid and a liquid, the rate of reaction is increased by

A. decreasing the surface area of the solid
B. increasing the size of the solid
C. increasing the surface area of the solid
D. reducing the concentration of the liquid

8.  Which of the following will slow down a chemical reaction?

A. Increasing temperature
B. Increasing concentration
C. Adding a catalyst
D. Lowering temperature

9. The rate of reaction between magnesium and dilute hydrochloric acid can be measured by

A. change in colour
B. volume of gas produced
C. mass of magnesium used
D. number of bubbles formed

10. Which graph best represents a fast reaction?

A. One that rises steeply
B. One that rises slowly
C. One that is flat
D. One that slopes downward

11. Which of the following graphs best represents an exothermic reaction?
A. Products higher than reactants
B. Products lower than reactants
C. Reactants and products at same level
D. No activation energy

12. The peak of an energy profile diagram represents the
A. reactants
B. products
C. activated complex
D. catalyst

13. The minimum energy required for a reaction to occur is called
A. enthalpy
B. heat of reaction
C. activation energy
D. reaction rate

14. In an endothermic reaction, the heat of reaction (ΔH) is
A. zero
B. negative
C. positive
D. constant

15. If a reaction produces a gas and the mass of the reaction mixture decreases with time, this is because
A. heat is absorbed
B. gas escapes
C. the solid melts
D. the reaction stops

16. A steep slope on a rate of reaction graph indicates
A. a slow reaction
B. no reaction
C. a fast reaction
D. equilibrium

17. Which of the following will increase the rate of a chemical reaction?
A. Decreasing temperature
B. Removing a catalyst
C. Increasing concentration
D. Decreasing surface area

18. In an exothermic reaction, the products are
A. more energetic than the reactants
B. less energetic than the reactants
C. equal in energy to the reactants
D. unstable

19. The heat energy absorbed or released during a reaction is represented by
A. Ea
B. ΔH
C. K
D. pH

20. On a volume of gas versus time graph, the reaction is complete when
A. the curve is steep
B. the curve is straight
C. the curve becomes horizontal
D. the volume decreases 

21. The rate of a reaction is given by the expression
{Rate} = k[A]^2
If the concentration of (A) is increased from 0.2 mol dm⁻³ to 0.4 mol dm⁻³, how does the rate of reaction change?

A. It doubles
B. It triples
C. It increases four times
D. It remains the same

22. A finely divided form of a metal burn more readily in air than the rod form because the rod has  (a) higher molar mass (b).  smaller surface area (c) protective oxide coating (d) different chemical properties 

THEORY QUESTIONS.

1.(a)(i) Define rate of a chemical reaction?

     ii). mention three factors that can affect the rate of a chemical reaction

     iii). state the collision theory

b).  A sample of carbon is burnt at a rate of 0.50gper second for 30 minutes to generate heat.

    (i). write a balanced equation for the reaction

   (ii). determine the    I. volume of carbon (IV) oxide produced at s.t.p.  II. moles of oxygen used up in the process at s.t.p. [ C= 12.0, O= 16.0, Molar volume Vm = 22.4dm³

c. State how each of the following affects the rate of chemical reactions (i) surface area (ii) catalysts [waec]

2.a(I) What is meant by the rate of chemical reaction? 

(ii). Explain in terms of collision theory, the effect of temperature increase on reaction rate.

(b) When hydrogen peroxide is exposed to air, it decomposes (I) write an equation for the reaction (ii). Outline an experiment to illustrate the effect of a named catalyst on the rate of decomposition. 

(iii).         

3a(i) Sketch an energy profile diagram to show the effect of a catalyst on the reaction rate, given that it is exothermic.

b. The graph below is the ratio curve for the following reaction carried out in an open vessel 

MgCO3(aq) + 2HCl → MgCl2(aq) +H2O(l) + CO2(g)

(i)  For how long the reaction occurred 

(ii) why was there a loss in mass?

(iii) State whether the reaction rate was fastest at the beginning, the middle  or towards the end of the reaction.  Give reason for your answer.

(iv) List three reaction conditions that can affect the slope of the curve.(waec)

4(a)In an experiment, excess 0.050mol/dm3 HCl was added to 10g of granulated zincin a beaker. Other conditions remaining constant state how the reaction rate would be affected in each case if the experiment was repeated using 

(i) 1.0mol/dm₃ HCl

(ii) 8.0g of granulated zinc

(iii) 10g of zinc dust 

(iv) a higher volume of 0.5mol/dm₃ HCl 

(v) a reaction vessel dipped in crushed ice 

(vi) equal volumes of water and 0.50mol/dm3 HCl.

b. 

5. (a) What is meant by the rate of a chemical reaction?

     (b) State three factors that affect the rate of a chemical reaction.

6. Explain how each of the following affects the rate of a chemical reaction:

(a) Temperature
(b) Concentration
(c) Surface area

7. (a) What is a catalyst?

(b) Give two examples of catalysts.
(c) State two characteristics of a catalyst.

8. A piece of calcium carbonate reacts with dilute hydrochloric acid.

(a) Write the balanced chemical equation for the reaction.
(b) State two ways in which the rate of the reaction can be increased.
(c) Explain one of the ways stated in (b).

9. (a) Explain why powdered zinc reacts faster with dilute hydrochloric acid than a lump of zinc.

 (b) What is meant by activation energy?

(c) How does a catalyst affect the activation energy of a reaction?

10. During a chemical reaction, the volume of gas produced was measured with time.

 (a) What information can be obtained from such a graph?
(b) How can you tell from the graph that the reaction is fast?

11. State two industrial processes where catalysts are used and name  the catalyst used in each case.

 (b) Explain why an increase in pressure increases the rate of reaction between gases.

(12).  (i) Draw the energy profile diagram for the reaction H2(g) + I2(g) → 2HI(g) △H = -13kJmol-1 if the concentration of HI(g) increases from 0.001 to 0.002 mol/dm3 in 80 secs what is the rate of the reaction