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Class 9 SELINA Solutions Chemistry Chapter 5 - The Periodic Table

The Periodic Table Exercise Ex. 5(A)

Solution 1

It is impossible for a chemist to study all the elements and their compounds. Hence, classification is a must.

Following are the reasons for the classification of elements:

  1. To study elements better
  2. To correlate the properties of the elements with some fundamental properties which are characteristic of all the elements
  3. To reveal relationships between elements 

Solution 2

The first classification of elements was into 2 groups-metals and non-metals.

Solution 3(a)

At. wt. of A = 7, At. wt. of C = 39

 

  

At. wt. of B = 23

i.e. Average of weights of A and C.

Solution 3(b)

  1. Döbereiner failed to arrange all the known elements in the form of triads.
  2. In the triad of fluorine (19), chlorine (35.5) and bromine (80), it is observed that the mean of the atomic masses of fluorine and bromine is ½(19 + 80) = 49.5, not 35.5. 

Solution 4

Elements when arranged in the increasing order of their atomic weights are similar to the eighth and the first note of the musical scale. For example, the eighth element from lithium is sodium. Similarly, the eighth element from sodium is potassium. Thus, lithium and sodium provide any specific place for hydrogen.

  1. This classification did not work with heavier elements.
  2. Newland adjusted two elements Cobalt (Co) and Nickel (Ni) in the same slot.
  3. Fe, which resembles Co and Ni in properties, has been placed far away. 

Solution 5

Yes, Döbereiner's triads also exist in the columns of Newland's octaves. For example, the second column of Newlands classification has the elements Lithium (Li), Sodium (Na) and Potassium (K), which constitute a Döbereiner's triad.

Solution 6(a)

Elements of lithium, sodium and potassium have the following similar properties:

  1. All these have one electron in the outermost shell.
  2. They form unipositive ions.
  3. They are good reducing agents.
  4. They are soft metals.
  5. They impart colour to the flame.
  6. Common name of the group is alkali metals [Group 1A]. 

Solution 6(b)

  1. All of them are metals.
  2. Oxide of each of them is alkaline in nature.
  3. Each has valency 2. 

Solution 7(a)

Mendeleev's basis for periodic classification:

  1. Similarities in the chemical properties of elements.
  2. Increasing order of atomic weights of elements. 

Solution 7(b)

Mendeleev laid the foundation for the modern periodic table by showing periodicity of the properties of the elements by arranging the elements (63) then known into 8 groups, by leaving gaps for undiscovered elements and predicting their properties. He made separate groups for metals and non-metals. He also created periods in which the element gradually changes from metallic to non-metallic character. He was also able to show that the element in the same sub-group had the same valency.

Solution 8

Mendeleev's periodic law: The physical and chemical properties of all the elements are a periodic function of their atomic masses.

Solution 9(a)

C is in Group 4. So, the hydride will be CH4 (Methane).

Si is in Group 4. So, the hydride will be SiH4 (Silane).

Solution 9(b)

K is in Group 1. So, the oxide will be K2O (Potassium oxide).

Al is in Group 3. So, the oxide will be Al2O3 (Aluminium oxide).

Ba is in Group 2. So, the oxide will be BaO (Barium oxide).

Solution 10

Anomalous pairs of elements were missing from Mendeleev's periodic table.

Solution 11

Merits of Mendeleev's classification of elements:

  1. Grouping of elements
  2. Gaps for undiscovered elements: Mendeleev left some gaps in his periodic table for subsequent inclusion of elements not known at that time.
  3. He predicted the properties of the then unknown elements on the basis of properties of elements lying adjacent to the vacant slots (eka-aluminium and eka-silicon). 

Solution 12

He left gaps in the table for the undiscovered elements. He discovered the properties of such elements with the help of neighboring elements.

He discovered eka-silicon with atomic mass of 72 which was later named Germanium with atomic mass 72.6.

Solution 13

Henry Moseley found that when cathode rays struck anodes of different metals, the wavelength of these metals was found to decrease in a regular manner of changing the metal of anode in the order of its position in the periodic table. By this, he concluded that the number of positive charges present in the nucleus due to protons (atomic number) is the most fundamental property of the element.

So, Henry Moseley found that the atomic number is a better fundamental property of an element compared to its atomic mass. This lead to the modern periodic law.

This law gave explanations for anomalies in Mendeleev's classification of elements such as

  1. Position of isotopes with the same atomic number can be put in one place in the same group.
  2. Position of argon and potassium: Potassium with higher atomic number should come later, and argon with lower atomic number should come first. 

Solution 14

 

 

Element

At. No.

Electronic distribution

Be

4

2, 2

Li

3

2, 1

Na

11

2, 8, 1

Ca

20

2, 8, 8

K

19

2, 8, 8, 2

 

  1. Same IA group (Li, Na, K) and IIA group (Be, Ca)
  2. In the second period (Be, Li) and in the fourth period (K, Ca) 

Solution 15(a)

Eka-silicon

Solution 15(b)

Gold and Platinum

Solution 15(c)

Only 63 elements were discovered at the time of Mendeleev's classification of elements.

The Periodic Table Exercise Ex. 5(B)

Solution 1(a)

Modern periodic law: The physical and chemical properties of all elements are a periodic function of their atomic numbers.

Solution 1(b)

Eighteen groups and seven periods

Solution 2

Last elements of each period have their outermost shell complete, i.e. 2 or 8 electrons.

The general name is inert gases or noble gases.

Solution 3(a)

Vertical columns in a periodic table which have the same number of valence electrons and similar chemical properties are called a group.

Solution 3(b)

In a periodic table, elements are arranged in the order of increasing atomic numbers in horizontal rows called periods.

Solution 4

Atomic number determines which element will be the first and which will be the last in a period of the periodic table.

Solution 5(a)

  1. Group 1 is known as the alkali metals.
  2. Group 17 is known as the halogens.
  3. Group 18 is known as the transition elements. 

Solution 5(b)

  1. Group 1: Lithium (Li), Sodium (Na)
  2. Group 17: Chlorine (Cl), Iodine (I)
  3. Group 18: Helium (He), Neon (Ne) 

 

Solution 6(a)

There are two elements in the first period.

Solution 6(b)

There are eight elements in the third period.

Solution 7(a)

  1. The valence electrons in the same shell (outermost shell) increase progressively by one across the period. The first element hydrogen has one valence electron and helium has two valence electrons.
  2. On moving from left to right in a period, valency increases from 1 to 4, then falls to one and ultimately to zero in the last group. 

Solution 7(b)

  1. Valence electrons in the same shell (outermost shell) increase progressively by one across the period. The first element sodium has one valence electron and magnesium has two valence electrons.
  2. On moving from left to right in the third period, valency increases from 1 to 7 and ultimately to zero in the last group. 

Solution 8

The size of atoms decreases when moving from left to right in a period. Thus, in a particular period, the alkali metal atoms are the largest and the halogen atoms are the smallest.

Li > Be > B > C > N > O > F

Solution 9(a)

  1. H and P are noble gases.
  2. G and O are halogens.
  3. A and I are alkali metals.
  4. D and L have valency 4. 

Solution 9(b)

Li2O. A stands for lithium and F stands for oxygen. The valence of lithium is +1 and the valence of O is -2, i.e. A2F.

Solution 9(c)

G has atomic number 9; therefore, its electronic arrangement is 2, 7.

Solution 10

Na and Al have the capacity to donate an electron due to which the valency is positive, whereas Cl and K can only gain or lose one electron due to which their valency is -1 and +1, respectively. This is the only difference between these two.

Solution 11

These elements have a full outermost subshell, which results in high stability. They only react with other elements in extreme circumstances.

Solution 12(a)

The greatest metallic character can be expected at the bottom of the group.

Solution 12(b)

The largest atomic size can be expected at the lower part of the group.

Solution 13

The number of valence electrons remains the same as we go down a group.

Solution 14(a)

Metals: A and B; Non-metals: C; Noble gases: D and E

Solution 14(b)

Most reactive

(i) Metals: Alkali metals (Group I); Caesium

(ii) Non-metals: Halogens (Group 17); Fluorine

Solution 14(c)

Element A will form a positive ion 1+ (cation).

Element B will form a positive ion 2+ (cation).

Element C will form a negative ion 1- (anion).

Solution 14(d)

  1. E
  2. B 

Solution 15

K L M

Electronic configuration = 2, 8, 7

  1. VIIA
  2. Third period
  3. Seven
  4. Valency of T = -1
  5. Non-metal
  6. Protons = 17, Neutrons = 18

 

The Periodic Table Exercise Ex. 5(C)

Solution 1

Atomic number of P = 19

Its electronic configuration = 2, 8, 8, 1

Group no. of the element = 1A

Period no. of the element = 4

P is a metal.

Solution 2

  1. 3
  2. +3
  3. Metal
  4. Aluminium 

Solution 3

  1. Helium
  2. Silicon
  3. 4, 3
  4. Argon
  5. Noble gases
  6. Carbon tetrachloride (CCl4)
  7. Silicon, Phosphorus
  8. Sodium chloride (Na+Cl-)
  9. Li and Mg; Be and Al; B and Si
  10. Sodium
  11. Typical elements of Period 2 belonging to Group 14 and 15 are carbon and nitrogen.

       Typical elements of Period 3 belonging to Group 14 to 15 are silicon and phosphorus.

  1. Beryllium

 

Solution 4

 

Column A

Answers

(a) Element short by 1 electron in octet

(v) Halogens

(b) Highly reactive metals

(iii) Alkali metals

(c) Non-reactive elements

(ii) Noble gases

(d) Elements of Groups 3 to 12

(i) Transition elements

(e) Radioactive elements

(vi) Actinides

(f) Elements with 2 electrons in the outermost orbit

(iv) Alkali earth metals

 

 

Solution 5

 

 

Atomic number

Element

Electronic configuration

Select element of the same group

11

Sodium

2, 8, 1

K

15

Phosphorus

2, 8, 5

N

16

Sulphur 

2, 8, 6

O

9

Fluorine

2, 7

Cl 

 

 

 

Solution 6

  1. Relative atomic mass of a light element up to calcium is approximately 20 its atomic number.
  2. The horizontal rows in a periodic table are called periods.
  3. Going across a period left to right, atomic size increases.
  4. Moving left to right in the second period, number of valence electrons increases from 1 to 8. 
  5. Moving down in the second group, number of valence electrons remain same.

Solution 7(a)

Name of the alkali metals: Lithium, sodium, potassium, rubidium, cesium and francium

Electrons in the outermost orbit: 1 

Solution 7(b)

  1. Reaction of alkali metal with oxygen - React rapidly with oxygen

        4Na + O2 → 2Na2O


  1. Reaction of alkali metal with water - React with water violently and produce hydrogen

       2M + 2H2O → 2MOH + H2


  1. Reaction of alkali metal with acid - React violently with dil. HCl and dil. H2SO4 to produce hydrogen

      2M + 2HCl → 2MCl + H2

 

Solution 8(a)

Alkali metals can be extracted by the electrolysis of their molten salts.

Solution 8(b)

The colour of the flame of sodium is golden yellow, and the colour of the flame of potassium is pale violet.

Solution 9

  1. Ca
  2. 1s22s22p63s23p64s2
  3. 2
  4. Group 2 Period 4
  5. Metal
  6. Reducing agent

 

Solution 10(a)

The first three alkaline earth metals are Beryllium, Magnesium and Calcium.

Solution 10(b)

Reactions of the first three alkaline earth metals with dilute hydrochloric acid:

Be + 2HCl → BeCl2 + H2

Mg + 2HCl → MgCl2 + H2

Ca + 2HCl → CaCl2 + H2

Solution 11(a)

Alkaline earth metals occur in nature in the combined state and not in the free state as they are very reactive.

Solution 11(b)

Electronic configuration of the first two alkaline earth metals:

4Be: 1s22s2

12Mg: 1s22s22p63s2

Solution 12

  1. Due to the reactive nature of alkali metals, they are kept in inert solvents.
  2. Alkali metals and halogens are very reactive; hence, they do not occur in the free state in nature.
  3. Alkali metals and alkaline earth metals have 1 and 2 valence electrons, respectively, in their outermost shell.
    They can lose electrons to atoms of non-metals to form an electrovalent compound.
  1. Inert gases have 2 or 8 electrons (duplet/octet) in their outermost orbit. That is their electronic arrangement is very stable, so they are unreactive and do not form compounds.

Solution 13

  1. Li < Na < K < Rb < Cs
  2. F > Cl > Br > I
  3. He < Na < Mg
  4. Cl < Mg < Na

Solution 14(a)

  1. Group 17 elements react with metals to form metal halides which are neutral in nature.
  2. Group 17 elements react with non-metals to form acidic compounds such as hydrogen halides. 

Solution 14(b)

Group 17 elements are highly reactive because of their closeness to the noble or stable gas configuration. They can easily achieve a noble gas electron structure.

Solution 15(a)

All the noble or inert gases have 8 electrons in their valence shell except helium which has two electrons in its valence shell.

Solution 15(b)

Xenon or krypton from Group 18 can form compounds.

Solution 16

  1. Helium
  2. Argon
  3. Neon

Solution 17(a)

Group 17 elements are called halogens. The name halogens is from Greek halo (sea salt) and gens (producing, forming) and thus means 'sea salt former'.

Solution 17(b)

Group 17 elements or halogens:

  1. Reactivity: Halogens are the most reactive non-metals, their reactivity decreases down the group. Fluorine is the most reactive halogen and iodine is the least reactive halogen.
  2. Colour: Fluorine is a pale yellow gas, chlorine is a greenish yellow gas, bromine is a reddish brown liquid and iodine is a violet solid.
  3. Physical state: Gaseous 

Solution 18

As elements P and Q belong to the same period of the modern periodic table and are in group 1 and group 2, they belong to alkali metals and alkaline earth metals, respectively.

Characteristic

Alkali metals (Element P)

Alkaline earth metals (Element Q)

(a) Number of electrons in their atoms

Valence electron = 1

Valence electrons = 2

(b) Their tendency to lose electrons

Can easily lose electrons

They will lose electrons easily but not as easily as alkali metals can.

(c) Formation of their oxides

React rapidly with oxygen in air

4Na + O2 → 2Na2O

Less reactive than alkali metals

(d) Formulae of their chlorides

NaCl, KCl, CsCl

MgCl2, CaCl2, BaCl2

 

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