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

     → PERIOD  

            ↓

         GROUP

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 (the highest positive integer) 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 element   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}

i.  They are known as oxygen family,

ii. They are all non -metals

iii. They ionize readily by gaining two electrons to form divalent negative ions.

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

v. They do not react with Water but oxygen combine directly with hydrogen to form water.

vi. They do not conductor electricity

vii. They are electro-negative   

vii. They electrons acceptor

 

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