ATOMIC NUMBER and ATOMIC MASS & RELATIVE ATOMIC MASS at a glance

CONSTITUENTS OF AN ATOM

The atom as explained earlier is made up of three sub-particles. 

i. Protons, 

ii. Electrons and 

iii. Neutrons. 

the Protons are positively charged. 

the Electrons are negatively charged    while 

the Neutrons have no charge.

S/n         sub-atomic particles        charge                mass

1.             Protons                           + (positive)            1  (unit)

2.             Electrons                         - (negative)            0.00005

3.             Neutrons                          0 (neutral)               1 (unit)                                   

Atomic number: - Of an element is the number of protons in the nucleus of the atom of the element.

atomic numbers are whole numbers and NO two elements in the world have the same atomic number. 

Mass number or atomic mass: - of an element is the sum of the protons and neutrons in the nucleus of its atom.

that is,

Atomic Mass = Number of protons + Number of neutrons

In chemistry an element X can be represented as                A

                                                                                              _ZX

where             A = Atomic mass or mass number

                      Z = Atomic number

Example      40                  mass no = 40 (20 protons + 20 neutrons)

                   20 Ca           atomic no = 20 (number of protons)

     Elements     Num. of  P        Num. of  E         Num. of  N

        40
1.    20 Ca                  20                    20                    20

   

     14
2.   7 N                       7                       7                    14

     

       39
3.    19 K                   19                   19                    20

 Examples 

1.            State the constituents of an atom.

2.            What is the number of protons, neutrons and electrons in the following elements

            11               12               23                  32
(a)        5B      (b)    6C    (c)   11Na     (d)   16S 

ISOTOPES

Isotopy is the occurrence of atoms of elements having the same atomic number but different mass numbers. This is due to the difference in the number of neutrons present in the atoms. The atoms that exhibit isotopy are called ISOTOPES.

Examples of atoms that exhibit isotopy are chlorine 35_Cl and 37_Cl

Carbon-  ¹² _C        ¹³ _C and ¹⁴ _C6

                              

Potassium – ³⁹K₁₉   and   ⁴¹K₁₉             

Oxygen -  ¹⁶O₁₆    and ¹⁸O₁₆

                                             

EVALUATION

1.            Define isotopy.

2.            Write the isotopes of chlorine.

                                             

CALCULATION OF RELATIVE ATOMIC MASS

RELATIVE ABUNDANCE:  this is the amount that a particular isotope occurs in nature, and it is expressed in percentage.

For example, Chlorine always occurs in two isotopic mixtures 35_Cl and 37_Cl with percentages of 75% and 25% respectively. This percentages are known as the relative abundance of each isotope

The following is an example of calculation of relative atomic mass of an element from percentage abundance of its isotopes.

1. X is an element which exists as an isotopic mixture containing 90% of ³⁹X₁₉ and 10% of ⁴¹X₁₉

a.    How many neutrons are present in the isotope ⁴¹X                                                                               

b.   Calculate the mean relative atomic mass of X

Solution

a.            Neutrons in  ⁴¹X_{19       :   mass number - atomic number} 

                                                         =          41-19 = 22

b.         R.A.M   =    90    x 39    +          10   x 41
                               100                          100

                                =          90  x   39  +   41  x  10
                                                          100

                                                         =  3920    =  39.20
                                                              100

1.         How many neutrons are present in the isotope ³⁷Cl₁₇ ?

2.         A given quantity of chlorine contains 75% ³⁵Cl₁₇, and 25% ³⁷Cl₁₇, determine the relative atomic mass of chlorine.

 

Example 2.

1. The following are more examples on calculations of relative atomic masses of elements.
2. An element Y exist in two isotopic forms ³⁹Y₁₈ and ⁴⁰Y₁₈ in the ratio 3:2 respectively. What is the relative atomic mass of the element?

SOLUTION                            

R.A.M of Y        =          3          x          39        +          2          x          40

                                    5                      1                     5                       1

                        =         0.6        x         39         +         0.4        x          40

                        =        23.4       +       16

                        =       39.4  

3.          An element with relative atomic mass 16.2 contains two isotopes 16P8 with relative abundance 90% and mP8 with relative abundance 10%. What is the value of m?

SOLUTION

            16.2 =  90 x 16 + 10 x m
                                  100

             16.2 = 9 x 16  +   m
                          10           10

            16.2 = 144   +   m
                        10         10

            16.2 = 144  +  m
                              10

            16.2 x 10 = 144 + m

            162   = 144 + m

            162 – 144 = m

                  18 =   m

The value of m is 18

OBJECTIVE QUESTIONS 

1.    The atomic number of an element is precisely.

(a) the number of protons in the atom

 (b) the number of electrons in the atom 

(c) the number of neutrons in the atom

(d)  the number of atoms in an element

2.    An atom can be defined more accurately as 

(a) the smallest indivisible parts of an element that can take part in a chemical reaction

 (b) the smallest part of an element that can take part in a chemical reaction

 (c) a combination of protons, neutrons

(d)  the smallest particle of a substance that can take part in a chemical reaction

3.      The mass number is 

(a) proton number + neutron number 

(b) electron number + proton number  

(c) neutron number + electron number

(d) atomic number + electron number

4.  Calculate the relative atomic mass of an element having two isotopes ¹⁰⁷ Ag and ¹⁰⁹Ag in the ratio 1:1  

(a)106             

(b)107         

(c)108     

(d) 109

5.   An element X has two isotopes ^18.8X and ^15.8X in the proportion of 1:9 respectively. Find the relative atomic mass of X           

 (a) 16.1    

(b) 13.6 

(c) 16.8 

 (d)  17

THEORY

1.            Consider the atoms represented below:

                                                                  ^q_X                      ^r_X
                                                        p                      p

a.            State the relationship between the two atoms.

b.            What is the difference between them?

c.             Give two examples of other elements which exhibit the phenomenon illustrated.  (waec)

2.            State the number of electrons, protons and neutrons present in the following atoms/ions

a)    Ca         b) S^2-         c) Al³⁺         d) P

b).            If an element R has isotopes 60% ¹²R₆ and 40% ^xR₆ and the relative atomic mass of R is 12.4, find x.

 3.(a)        Define the term isotopy.

(b)        Determine the number of electrons, protons and neutrons in each of the following: ³⁹K₁₉, ^63.5Cu₂₉

c)         If an element R has isotopes 60% of ¹²R₆ and 40%  ^xR₆ and the relative atomic mass is 12.4, find x.

     

4.        How many neutrons are present in the isotope ³⁷Cl₁₇ ?

5.        A given quantity of chlorine contains 75% ³⁵Cl₁₇, and 25% ³⁷Cl₁₇, determine the relative atomic mass of chlorine.

 

QUANTUM NUMBERS at a glance

 

Rules guiding the arrangement of electrons in an atom

When electrons are arranged or filled into the atoms of elements, certain RULES are considered   and obeyed. These rules are the Aufbau's Principle, Pauli's exclusion principle and Hund’s rule of maximum multiplicity.

PAULI EXCLUSION PRINCIPLE 

states that NO two electrons in the same atom have the sets of the four quantum numbers {n, l, m and s in an atom}.

AUFBAU PRINCIPLE states that "when electrons go into atoms they fill orbitals of lower energy first before filling orbitals of higher energy and each orbital may hold up to two electrons.

 

HUND’S RULE OF MAXIMUM MULTIPLICITY state that " When electrons fill degenerate orbitals they go in singly first before pairing up occurs.

 Degenerate Orbitals are orbitals that are at the same energy level. example of degenerate orbitals is the P-orbital, the d-orbital or the f-orbital

Examples of degenerate orbitals are the P-orbitals,  the d-orbitals  and the f-orbitals 

QUANTUM NUMBERS

The quantum numbers are a set of numbers that describes the position of an electron in an atom. 

Studies have showed  that the energy of an electron may be characterized by four quantum numbers. These quantum numbers help to locate the position of electrons in an atom

1. The principal quantum number represented by n with integral values of 1,2,3,4 e.t.c.

This quantum number describes the shell ( k, l, m....)

2. The subsidiary or Azimuthal quantum number represented by l with integral values ranging from 0 to (n-1).

This quantum number describes the sub-shells ( s,p,d,f,)

3.  The magnetic quantum number represented by m with integral values ranging from –l ,0, +l.

4. The spin quantum number represented by s with integral values – ¹/₂ and + ¹/₂.

Element   At. Numb.  Elect. Conf.

 H .            1;             1s¹

He              2;         1s²

Li.               3;         1s² 2s¹

Be              4;          1s² 2s²

B =.           5;          1s² 2s² 2p¹

C =            6;          1s² 2s² 2p²

N =.          7 ;         1s²2s²2p3      

O=            8 ;       1s² 2s² 2p⁴

F=.           9;         1s² 2s² 2p⁵

Ne=        10;         1s² 2s² 2p⁶

Na=.         11;        1s² 2s² 2p⁶ 3s¹

Mg=          12;     1s² 2s² 2p⁶ 3s²

Al=.        13;   1s² 2s² 2p⁶ 3s² 3p¹

Si =       14;    1s² 2s² 2p⁶ 3s² 3p²

P=.       15;     1s² 2s² 2p⁶3s²³3p³

S =.       16;   1s² 2² 2p⁶3s² 3p⁴

Cl =.    17;     1s² 2s² 2p⁶ 3s² 3p⁵

Ar =.   18      1s² 2s² 2p⁶ 3s²3p⁶

K=.    19  1s² 2s² 2p⁶ 3s²3p⁶ 4s¹

Ca =.  20  1s² 2s² 2p⁶ 3s²3p⁶ 4s²

OBJECTIV QUESTION

1.  Which of the following orbitals is spherical in shape?

     (a) s

      (b) p

      (c) d 

      (d) f

  

2.     Which of the following shells have a maximum of eight electrons?

             (a)  K

             (b) L 

             (c) M 

             (d) N

3.     1s² 2s² 2p⁶ 3p¹ is the electronic configuration of

              (a)  potassium

               (b) calcium 

               (c) sodium

              (d) aluminum.

4  . “Two electrons in an atom cannot have the same set for all four quantum numbers”. This statement is

              (a)  Aufbau principle 

               (b) Pauli exclusion 

               (c) Hund’s rule

              (d) Rutherford’s model.

5.    Which of the quantum number is represented by L?

              (a)  principal quantum number

              (b) subsidiary quantum number                  

              (c) magnetic quantum number

              (d) spin quantum.

6.    Pauli exclusion principles related 

a). quantity of electrons in the valence shell

b). filling the orbitals with lower energy first 

c). the filling of degenerate orbitals 

d). quantum numbers of electrons.

7. Atomic orbital is 

a). the circular path through which electrons which electrons revolve round the nucleus 

b). a region around the nucleus where electrons are most likely to be found 

c). the path around the nucleus through which electrons move 

d). the path around the nucleus through which protons move.

THEORY QUESTIONS 

  1(a)(i)what are the quantum numbers 

   (ii). The models below represent the filling of orbitals in an atom

   

State which rule(s) is/are violated or obeyed by each model

(iii).     State the following principle (a) Pauli exclusion principle. (b) Aufbau principle

(b). Write the electronic configuration of 

(i) Oxygen   (ii) Calcium (iii) Fluoride ion (Cl^-) (iv) Potassium ion  (K⁺)    (v). Aluminum ion (Al³⁺

2.a(i). List the quantum numbers that are assigned to an electron in an atom.

(ii) what is the maximum number of electrons that can occupy the 3d orbital?

3.(a)State 

I.  Pauli's exclusion principle 

ii. Hund's rule of maximum multiplicity

iii. Write the electronic configuration of of each of the following ions of copper I. Cu+ 

II. Cu2+.       [29Cu]

THE HALOGENS at a glance

 

HALOGENS

Halogens (salt formers) are found in group VII of the periodic table. The halogens  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.

 In the group 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.

The periodic law states that the elements on the periodic table are arranged in order of their atomic number. OR The arrangements of the elements on the periodic table is a function of their atomic   number.

  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

Each horizontal row is called a period

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 represents the period to which the element belongs to while the number of electrons in the outermost shell of the configuration represents the group to which the element belongs. For example, Given two elements X and Y with the following electronic configuration X=1s²2s²2p⁴ and another element   Y = 1s²2s²2p⁶3s². The largest number in X is 2 (i.e X contains 2 shells) and hence belongs to period 2. The largest number in Y is 3 (i.e Y contains 3 shells) and belongs to period 3. The total number of electrons in the outermost shell of X is 6 (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 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 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 on 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.

IONIC RADIUS: -For metals their atomic radius is larger than their ionic radius, 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.

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.

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)

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.

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.

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

They are soft, malleable, and ductile

They ionize by loss of one electron

They are good reducing agents

They are good conductors of heat and electricity

They react with cold water to displace hydrogen gas

Na(s) + H2O(l) → NaOH(aq) + H2(g)

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

They ionize by the loss of two electrons 

They are good conductors of heat and electricity

They are good reducing agents

GROUP VI

 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} 

^{I. They} 

GROUP VII: - (Halogens)

They ionize by the gain of one electron

They are good oxidizing agents

They are coloured

-Florine is yellowish 

-chlorine is greenish yellow 

- bromine is reddish-brown 

-iodine is violet

GROUP VIII (Noble gases) (rare gases) (inert gases)

They are unreactive

They have complete octet or duplet structure

TRANSITION METALS: (d-block elements)

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

  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

OBJECTIVE QUESTIONS 

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

1. Which of the letters indicate elementswhich exist as diatomic gases.

a).  B and G

b).  Cand  F

c).  Cand 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).  Cand 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 

P). 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). Cand 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 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 elements of atoms 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 --- 1s²2s² 

i.  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 

ISOMERISM at a glance

 

ISOMERISM IN ORGANIC COMPOUNDS

Isomerism is defined as the occurrence of two or more compounds with the same molecular formular but different molecular structures. 

The different molecular structures are known as Isomers.

Isomers may belong to the same homologous series ( having the same functional group) or may belong to different homologous series (having different functional group)

Isomers with the same functional group, (i.e. belonging to the same homologous series) have similar chemical properties while isomers having different functional group (i.e. belonging to different homologous series) have different physical and chemical properties.

Example:- the formula C2H6O has two isomers 

1. Ethanol  an Alkanol with molecular formula  CH3CH2OH has a melting point of 78 degree celçius and is a liquid at room twmperature 

     H   H

H--C--C--OH

     H. H

and 

2. Methoxymethane an Alkoxy compound with molecular formula CH3OCH3 which is a gaseous at room temperature

     H       H

H--C--O--C--H

     H       H

TYPES OF ISOMERISM

(i) Structural isomerism

(ii) Stereo or geometric Isomerism

(iii) Optical isomerism

 

(I) STRUCTURAL ISOMERISM

This is the occurrence of two or more compounds with the same molecular formular but different molecular structures 

TYPES OF STRUCTURAL ISOMERISM

(a) Chain isomerism: in this case, the isomer differs in the way the carbon atoms are arranged in the molecule. Eg

 CH₃CH₂CH₂CH₃           CH₃CHCH₃    
 n-butane                                 |                                                                    
                                               CH₃   
                             2-methylpropane (isobutane)

(‘n’ stands for the normal compound while ‘iso’ means the isomer of the normal compound).

NOTE: - Isomerism begins from the 4th member of the alkane homologous series, that is, butane, and as the number of carbon atoms increases the number of isomers also increases. For example, butane has two isomers normal-butane and 2-methyl propane, pentane (C₅H₁₂) has three isomers:  n-pentane, 2-methylbutane and 2,2-dimethylpropane.            

  I.  CH₃CH₂CH₂CH₂CH3.  (n-pentane)

              CH₃
               |
 II.  CH₃CHCH₂CH3  (2-methylbutane)

                                     

III.              CH₃
                    |         
          CH₃-C-CH₃
                    |
                   CH₃        2,2-dimethylpropane

Three isomers for pentane 

Another example is hexane (C6H14)

I.  CH₃CH₂CH₂CH₂CH2CH3

               n-hexane

          

              CH₃
               |
 II.  CH₃CHCH₂CH2CH3    2-methylpentane

III.                        CH₃
                              |
              CH₃CH2CHCH2CH3          3-methylpentane

                   CH₃
                       |
 IV.      CH₃CCH₂CH3    2,2-dimethylbutane
                   |
                  CH3

                                     

                  CH₃
                    |
 V.      CH₃CHCHCH3    2,3-dimethylbutane
                        |
                       CH3  

Hexane has 5 isomers 

(b) position isomerism: Position isomers are those that have a substituent in different position on the same carbon skeleton. Example

                CH₃CH₂CH₂OH                 CH₃CHCH₃ 
                    propan-1-ol                             |
                                                                    OH 
                                                             propan-2-ol

(c) Functional Group isomerism: These are isomers having the same molecular formular but different functional groups. They belong to different homologous series.

e.g.
                         H
                        /
(i)  CH₃CH₂C                          CH₃-CCH₃ 
                       \\                                 ‖‖    
                        O                               O
     Propanal  (an aldehyde)       propan-2-one (a ketone)

(ii) CH₃CH₂OH                                  CH₃-O-CH₃ 

      Ethanol  ( Alkanol)                Methoxymethane (dimethyl ether) (Alkoxy compounds)

(2) STEREO ISOMERISM OR GEOMETRIC ISOMERISM

In this type of isomerism, compounds have the same molecular formular but differ only in the way their atoms are arranged in space.

TYPES   STEREO ISOMERISM

(a) Geometrical isomerism: This type of isomerism is found in compounds having either a double bond, triple bond or a ring structure. These multiple bonds prevent the free rotation about a carbon=carbon atom. that is, the carbon atoms are not cylindrically symmetrical. There are two forms of geometric isomers, the Cis and Trans isomer. For example, but-2-ene has two geometric isomer which are trans-but-2-ene and cis-but-2-ene as shown below.

when naming geometric isomers, you look at the substituents, where the substituents (group other than Carbon atoms) are on the same side of the double bond, that isomer is named the Cis isomer and when the substituents are on opposite sides of the double bond, we put a trans- isomer, but the names are the same

   H               CH₃                       H             H          
        \           /                                 \           /
           C=C                                       C=C
         /         \                                    ∕         ∖
  CH₃          H                            CH₃         CH₃

Trans-but-2-ene                          Cis-but-2-ene

 

  H                Br               H             H
      \               /                    \          /
          C=C                            C=C
      /             \                      /          \   
  Br              H                 Br          Br
Trans-1,2-dibromoethene       Cis-1,2-dibromoethene                    

   C1OOH.  C4OOH            H           COOH
     \           /                             \         /       
        C2=C3                              C=C    
       ∕         ∖                            /        \
COOH       H                      H             H
Trans-butenedioic acid            Cis-butenedioic acid

(b) Optical isomerism: Optical isomers   are isomers that have the same molecular and structural formular but cannot be superimposed on each other. In other words, an optical isomer is one which is not super-imposable on its mirror image. An optical isomer has at least one carbon atom which has four different groups or atoms attached to it. Such a carbon atom (surrounded by four different atoms /group) is called chiral carbon or chiral center. E.g. 2-hydroxylpropanoic acid (Lactic acid).

                           CH₃
                             |
              HO — *C— H
                             |
                            COOH

                  _{The Asterix C-atom is called a chiral C-atom because it is surrounded by four (4) different groups.}

            In glucose, there are several chiral carbon atoms.

                                          H
                                           ∣
                                          C = O
                                           |
                                H — *C —OH
                                            |
                              OH —*C —H
                                            |
                                H — *C— OH
                                           |
                                 H —*C —OH
                                           |
                                 H — C—H
                                           |
                                          OH

Optical isomers are also called Enantiomers.

Isomers that belong to the same homologous series (same functional group) have similar chemical properties but may have different physical properties. But isomers which belong to different homologous series (different functional group) have different chemical and physical properties.