Chapter 3 Classification of Elements and Periodicity in Properties- Study Materials

Subtopics of Class 11 Chemistry Chapter 3 Classification of Elements and Periodicity in Properties

  1. Why Do We Need To Classify Elements?
  2. Genesis Of Periodic Classification
  3. Modern Periodic Law And The Present Form Of The Periodic Table
  4. Nomenclature Of Elements With Atomic Numbers > 100
  5. Electronic Configurations Of Elements And The Periodic Table
  6. Electronic Configurations And Types Of Elements: 5-, P-, D-, F- Blocks Ex
    • The S-block Elements
    • The P-block Elements
    • The D-block Elements (Transition Elements)
    • The F-block Elements (Inner-transition Elements)
    • Metals, Non-metals And Metalloids
  7. Periodic Trends In Properties Of Elements
    • Trends in Physical Properties
    • Periodic Trends In Chemical Properties
    • Periodic Trends And Chemical Reactivity

Classification of Elements and Periodicity in Properties Class 11 Notes Chemistry Chapter 3

• Genesis of Periodic Classification Dobereiner’s Triads
In 1829, Dobereiner arranged certain elements with similar properties in groups of three in such a way that the atomic mass of the middle element was nearly the same as the average atomic masses of the first and the third elements. A few triads proposed by him are listed.
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Limitations of Dobereiner’s Trids
The triads given by Dobereiner were helpful in grouping some elements with similar characteristics together, but he could not arrange all the elements known at that time into triads.
• Newlands’ Law of Octaves
John Newlands proposed the law of octaves by stating that when elements are arranged in order of increasing atomic masses, every eighth element has properties similar to the first. Newlands called it law of octaves because similar relationship exists in the musical notes also.
This can be illustrated as:
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Limitations of Newlands’ Law of Octaves
(i) This classification was successful only up to the element calcium. After that, every eighth element did not possess the same properties as the element lying above it in the same group.
(ii) When noble gas elements were discovered at a later stage, their inclusion in the table disturbed the entire arrangement.
• Mendeleev’s Periodic Table
Mendeleev’s Periodic Law: The physical and chemical properties of the elements are a periodic
function of their atomic masses.
Mendeleev arranged the elements known at that time in order of increasing atomic masses
and this arrangement was called periodic table.
Elements with similar characteristics were present in vertical rows called groups. The horizontal
rows were known as periods.
Description of Mendeleev’s Periodic Table
(i) In the periodic table, the elements are arranged in vertical rows called groups and horizontal rows known as periods.
(ii) There are nine groups indicated by Roman Numerals as I, II, III, IV, V, VI, VII, VIII and zero. Group VIII consists of nine elements which are arranged in three triads. The zero group contains elements belonging to inert gases or noble gases and elements present have zero valency.
(iii) There are seven periods (numbered from 1 to 7) or, horizontal rows in the Mendeleev’s periodic table.
Importance of Mendeleev’s Periodic Table
(i) This made the study of the elements quite systematic in the sense that if the properties of one element in a particular group are known, those of others can be pridicted.
(ii) This helped to a great extent in the discovery of these elements at a later stage.
(iii) Mendeleev corrected the atomic masses of certain elements with the help of their expected positions and properties.
Defects in Mendeleev’s Periodic Table
(i) Hydrogen has been placed in group IA along with alkali metals. But it also resembles halogens of group VII A in many properties. Thus, its position is the Mendeleev’s periodic table is controversial.
(ii) Although the elements in the Mendeleev’s periodic table have been arranged in order of their atomic masses, but in some cases the element with higher atomic mass precedes the element with lower atomic mass.
(iii) We know that the isotopes of an element have different atomic masses but same atomic number. Since, periodic table has been framed on the basis of increasing atomic masses of the elements, different positions must have been allotted to all the isotopes of a particular element.
(iv) According to Mendeleev, the elements placed in the same group must resemble in their properties. But there is no similarity among the elements in the two sub-groups of a particular group.
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(v) In some cases, elements with similar properties have been placed in different groups.
(vi) Lanthanoids and actinoids were placed in two separate rows at the bottom of the periodic table without assigning a proper reason.
(vii) No proper explanation has been offered for the fact that why the elements placed in group show resemblance in their properties.
• Modern Periodic Law
Physical and chemical properties of the elements are the periodic function of their atomic numbers.
• Present Form of the Periodic Table (Long form of Periodic Table)
The long form of periodic table, also called Modem Periodic Table, is based on Modern periodic law. In this table, the elements have been arranged in order of increasing atomic numbers.
• Nomenclature of Elements with Atomic No. more than 100
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• Structural Features of the Periodic Table
Groups
The long form of periodic table also consists of the vertical rows called groups. There are in all 18 groups in the periodic table. Unlike Mendeleev periodic table, each group is an independent group.
Characteristics of groups:
(i) All the elements present in a group have same general electronic configuration of the atoms.
(ii) The elements in a group are separated by definite gaps of atomic numbers (2, 8, 8,18, 18,32).
(iii) The atomic sizes of the elements in group increase down the group due to increase the number of shells.
(iv) The physical properties of the elements such as m.p., b.p. density, solubility etc., follow a systematic pattern.
(v) The elements in each group have generally similar chemical properties.
Periods
Horizontal rows in a periodic table are known as periods.
There are in all seven periods in the long form of periodic table.
Characteristics of periods:
(i) In all the elements present in a period, the electrons are filled in the same valence shell.
(ii) The atomic sizes generally decrease from left to right.
s-Block Elements
General electronic configuration: ns1-2 Characteristics of s-block elements:
(i) All the elements are soft metals.
(ii) They have low melting and boiling points.
(iii) They are highly reactive.
(iv) Most of them impart colours to the flame.
(v) They generally form ionic compounds.
(vi) They are good conductors of heat and electricity. p-Block Elements
General electronic configuration: ns2np1-6
Characteristics of p-block elements:
(i) The compounds of these elements are mostly covalent in nature.
(ii) They show variable oxidation states.
(iii) In moving from left to right in a period, the non-metallic character of the elements increases.
(iv) The reactivity of elements in a group generally decreases downwards.
(v) At the end of each period is a noble gas element with a closed valence shell ns2 np6 configuration.
(vi) Metallic character increases as we go down the group.
d-Block Elements
General electronic configuration: (n -1) d1-10 ns0-2
The d-block elements are known as transition elements because they have incompletely filled d-orbitals in their ground state or in any of the oxidation states.
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Characteristics of d-block elements:
(i) They are all metals with high melting and boiling points.
(ii) The compounds of the elements are generally paramagnetic in nature.
(iii) They mostly form coloured ions, exhibit variable valence (oxidation states).
(iv) They are of tenly used as catalysts.
f-Block Elements
General electronic configuration: (n – 2) f1-14 (n -1) d0-1 ns2
They are known as inner transition elements because in the transition elements of d-block, the electrons are filled in (n – 1) d sub-shell while in the inner transition elements of f-block the filling of electrons takes place in (n – 2) f subshell, which happens to be one inner subshell. Characteristics of f-Block elements:
(i) The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids Ce (Z = 58) – Lu (Z = 71) and Actinoids Th (Z = 90) – Lr (Z = 103).
(ii) These two series of elements are called Inner Transition Elements (f-Block Elements).
(iii) They are all metals. Within each series, the properties of the elements are quite similar.
(iv) Most of the elements pf the actinoid series are radio-active in nature.
• Metals
(i) Metals comprise more than 78% of all known elements and appear on the left side of the Periodic Table.
(ii) Metals are solids at room temperature.
(iii) Metal usually have high melting and boiling points.
(iv) They are good conductors of heat and electricity.
(u) They are malleable and ductile.
• Non-metals
(i) Non-metals are located at the top right hand side of the Periodic Table.
(ii) Non-metals are usually solids or gases at low temperature with low melting and boiling points.
(iii) They are poor conductors of heat and electricity.
(iv) The non-metallic character increases as one goes from left to right across the Periodic
Table.
(v) Most non-metallic solids are brittle and are neither malleable nor ductile.
• Metalloids
The elements (e.g., silicon, germanium, arsenic, antimony and tellurium) show the characteristic, of both metals and non-metals. These elements are also called semimetal.
• Noble Gases
– These are the elements present in group 18.
– Eash period ends with noble gas element.
– All the members are of gaseous nature and because of the presence of all the occupied filled orbitals, they have very little tendency to take part in chemical combination.
– These are also called inert gases.
• Representative Elements
The elements of group 1 (alkali metals), group 2 (alkaline earth metals) and group 13 to 17 constitute the representative elements. They are elements of s-block and p-block.
• Transition Elements
The transition elements include, all the d-block elements and they are present in the centre of the periodic table between s and p-block elements.
• Inner Transition Elements
Lanthanoids (the fourteen elements after Lanthanum) and actinides (the fourteen elements after actinium) are called inner transition elements. They are also called f-block elements.
The elements after uranium are also called transuranic elements.
• Periodic Trends in Properties of Elements
Trends in Physical Properties
Atomic Radii: It is defined as the distance from the centre of the nucleus to the outermost shell containing the electrons. Depending upon whether an element is a non-metal or a metal, three different types of atomic radii are used. These are:
(a) Covalent radius (b) Ionic Radius (c) van der Waals radius (d) Metallic radius.
(a) Covalent Radius: It is equal to half of the distance between the centres of the nuclei of two atoms held together by a purely covalent single bond.
(b) Ionic Radius: It may be defined as the effictive distance from the nucleus of an ion upto which it has an influence in the ionic bond.
(c) van der Waals Radius: Atoms of Noble gases are held together by weak van der Waals forces of attraction. The van der Waals radius is half of the distance between the centre of nuclei of atoms of noble gases.
(d) Metallic Radius: It is defined as half of the intemuclear distance between the two adjacent metal ions in the metallic lattice.
• Variation of Atomic Radius in the Periodic Table
Variation in a Period: Along a period, the atomic radii of the elements generally decreases from left to right.
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Variation in a group: The atomic radii of the elements in every group of the periodic table increases as we move downwards.
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• Ionic Radius
The ionic radii can be estimated by measuring the distances between cations and anions in ionic crystals.
In general, the ionic radii of elements exhibit the same trend as the atomic radii.
Cation: The removal of an electron from an atom results in the formation of a cation. The radius of cation is always smaller than that of the atom.
Anion: Gain of an electron leads to an anion. The radius of the anion is always larger than that ‘ of the atom.
Isoelectronic Species: Some atoms and ions which contain the same number of electrons, we call them isoelectronic species. For example, O2-, F, Na+ and Mg2+ have the same number of electrons (10). Their radii would be different because of their different nuclear charges.
• Ionization Enthalpy
It is the energy required to remove an electron from an isolated gaseous atom in its ground state.
M (g) + I.E ——->M+ (g) + e
The unit of ionization enthalpy is kJ mol-1 and the unit of ionization potential is electron volt per atom.
Successive Ionization Enthalpies
If a gaseous atom is to lose more than one electron, they can be removed one after the other i.e., in succession and not simultaneously. This is known as successive ionization enthalpy (or potential).
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• Variation of Ionization Enthalpies in the Periodic Table:
Variation of Ionization Enthalpy Along a Period
Along a period ionization enthalpies are expected to increase in moving across from left to the right, because the nuclear charge increases and the atomic size decreases.
Variation of Ionization Ethalpy in a Group
The ionization enthalpies of the elements decrease on moving from top to the bottom in any group.
The decrease in ionization enthalpies down any group is because of the following factors.
(i) There is an increase in the number of the main energy shells (n) in moving from one element to the other.
(ii) There is also an increase in the magnitude of the screening effect due to the gradual increase in the number of inner electrons.
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• Electron Gain Enthalpy
Electron Gain Enthalpy is the energy released when an electron is added to an isolated gaseous atom so as to convert it into a negative ion. The process is represented as:
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For majority of the elements the electron gain enthalpy is negative. For example, the electron gain enthalpy for halogens is highly negative because they can acquire the nearest noble gas configuration by accepting an extra electron.
In contrast, noble gases have large positive electron gain enthalpies because the extra electron has to be placed in the next higher principal quantum energy level thereby producing highly unstable electronic configuration.
Successive Electron Gain Enthalpies
We have studied that electrons from a gaseous atoms are lost in succession (i.e., one after the other). Similarly, these are also accepted one after the other, i.e., in succession. After the addition of one electron, the atom becomes negatively charged and the second electron is to be added to a negatively charged ion. But the addition of second electron is opposed by electrostatic repulsion and hence the energy has to be supplied for the addition of second electron. Thus the second electron gain enthalpy of an element is positive.
For example, when an electron is added to oxygen atom to form O ion, energy is released. But when another electron is added to 0- ion to form O2- ion, energy is absorbed to overcome the strong electrostatic repulsion between the negatively charged 0 ion and the second electron being added. Thus, first electron gain enthalpy:
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Factors on which Electron Gain Enthalpy Depends
(i) Atomic size: As the size of an atom increases, the distance between its nucleus and the incoming electron also increases and electron gain enthalpy becomes less negative,
(ii) Nuclear charge: With the increase in nuclear charge, force of attraction between the nucleus and the incoming electron increases and thus electron gain enthalpy becomes more negative.
(iii) Symmetry of the Electronic Configuration: The atoms with symmetrical configuration (having fully filled or half filled orbitals in the same sub-shell) do not have any urge to take up extra electrons because their configuration will become unstable.
In that case the energy will be needed and electron gain enthalpy (Δ eg H) will be positive. For example, noble gas elements have positive electron gain enthalpies.
Variation of Electron Gain Enthalpy Across a Period
Electron gain ethalpy becomes more negative with increase in the atomic number across a period.
Variation of Electron Gain Enthalpy in a Group
Electron gain enthalpy becomes less negative as we go down a group.
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• Electronegativity 
A qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself is called electronegativity. Unlike ionization enthalpy and electron gain enthalpy, it is not a measurable quantity.
However, a number of numerical scales of electronegativity of elements viz, Pauling scale, Milliken- Jaffe scale, Allred Kochow scale have been developed. The electronegativity of any given element is not constant; it varies depending on the element to which it is bound.
Across a Period
Electronegativity generally increases across a period from left to right.
In a Group
It decreases down a group.
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• Periodic Trends in Chemical Properties along a Period
(i) Metallic character: Decrease across a period maximum on the extreme left (alkali metals).
(ii) Non-metallic character: Increasess along a period. (From left to right).
(iii) Basic nature of oxides: Decreases from left to right in a period.
(iv) Acidic nature of oxides: Increases from left to right in a period.
• Variation from Top to Bottom on Moving Down a Group
(i) Metallic character. Generally increases because increase in atomic size and hence decrease in the ionizatiort energy of the elements in a group from top to bottom.
(ii) Non-metallic character. Generally decreases down a group. As electronegativity of elements decreases from top to bottom in a group.
(iii) Basic nature of oxides. Since metallic character or electropositivity of elements increases in going from top to bottom in a group basic nature of oxidise naturally increases.
(iv) Acidic character of oxides. Generally decreases as non-metallic character of elements decreases in going from top to bottom in a group.
(v) Reactivity of metals. Generally increases down a group. Since tendency to lose electron increases.
(vi) Reactivity of non-metals. Generally decreases down the group, Higher the electro-negativity of non-metals, greater is their reactivity. Since electronegativity of non-metals in a group decreases from top to bottom, their reactivity also decreases.
• Anomalous Properties of Second Period Elements
The first element of each of the group 1 (lithium) and 2 (beryllium) and group 13-17 (boron to fluorine) differs in many respect from the other members of their respective groups. For example, lithium unlike other alkali metals, and beryllium unlike other alkaline earth metals
form compounds which have significant covalent character; the other members of these groups, pre-dominatly form ionic compounds.
It has been observed that some elements of the second period show similarities with the elements of the third period placed diagonally to each other, though belonging to different groups.
For example,
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This similarity in properties of elements placed diagonally to each other is called diagonal relationship.
• Mendeleev’s Periodic Law. Physical and chemical properties of elements are periodic function of their atomic masses.
• Modem Periodic Law. Physical and chemical properties of the elements are periodic function of their atomic numbers.
• Groups. There are 18 groups. These are vertical rows.
• Periods. There are 7 periods. These are horizontal rows.
• Representative Elements. The S and P block of elements are known as representative elements.
• Transition Elements. They are also called d-block elements. They have general electronic configuration (n – 1) d1-10 ns0-2.
• Inner Transition Elements. Lanthanoids (the fourteen elements after Lanthanum) and actinides (the fourteen elements after actinium) are called inner transition elements. General electronic configuration is (n – 2) f1-14(n – 1) d0-1 ns2.
They are also called f-block elements.
• Metals. Present on the left side of the periodic table. Comprise more than 78% of the known elements.
• Non-metals. Mostly located on the right hand side of the periodic table.
• Metalloids. Elements which line as the border line between metals and non-metals (e.g., Si, Ge, As) are called metalloids or semimetals.
• Atomic Radii and Ionic Radii, increase down the group decrease along the period.
• Ionization Enthalpy. Increases along the period and decreases down the group.
• Noble Gas Elements. Elements with symmetrical configuration are chemically inert in nature.
• Electric Nuclear Charge. Z = Nuclear charge – Screening constant.
• Electronegativity. Increases along a period decreases down the group,
• Chemical Reactivity. Chemical reactivity is highest at the two extremess of a period and lowest in the centre.
• Oxides of Elements. Oxides formed of the Elements on the left are basic and of elements
on the right are acidic in nature.
Oxides of elements in the centre are amphoteric or neutral.

CBSE Class 11 Chemistry Chapter-3 Important Questions


1 Marks Questions

1. How many elements are known at present? 

Ans.There are about 114 elements known at present.


2. Who was the first scientist to classify elements according to their properties?

Ans.The German Chemist, Johann Dobereiner in early 1829 was the first to consider the idea of trends among properties of element.


3. What is the basis of triad formation of elements?

Ans. The middle element of each of the triads had an atomic weight about half way between the atomic weights of the other two. Also the properties of the middle element were in between those of the other two members. Dobereiner’s relationship is known as the law of triads.


4. Stale the modern ‘Periodic law’?

Ans. The physical and chemical properties of the elements are periodic functions of their atomic numbers.


5. Define and state Mendeleev’s periodic law.

Ans.  Mendeleev’s Periodic law states that

‘The properties of the elements are periodic function of their atomic weights’.


6. Give the general characteristics of the long form of Modern periodic table?

Ans.  General characteristics of the long form of Periodic table :-

(i)There are in all 18 vertical columns i.e. 18 groups in the long form periodic table.

(ii)There are groups numbered from 1 to 18 from the left.

(iii)There are seven horizontal rows called periods.

(iv)The elements of groups 1, 2 and 13 to 17 are called main group elements.

(v)The elements of group 3 to 12 are called transition elements.


7. In short give the features of the seven periods. 

Ans. First period contains 2 elements, 1H and 2 He and it is the shortest period.

Second and third periods contain 8 elements each namely Li to 10Ne and 11Na to 18Ar and is a short period.

Fourth and fifth period contains 18 elements each namely 19K to 36Kr and 37Rb to 54Xe and is a long period.

Sixth period contains 32 elements namely 55Cs to 86Rn and is the longest period.

Seventh period is incomplete. It has all other elements starting with 87Fr onwards. Elements from 93 onwards are purely synthetic and are called trans–uranium elements and is incomplete period.


8. Define electronic configuration.

Ans.  The distribution of electrons into orbitals of an atom is electronic configuration.


9. What is the electronic configuration when elements are classified group wise?           

Ans. Elements in the same vertical column or group have similar valence shell electronic configurations, the same number of electrons in the outer orbitals, and similar properties.


10. Predict the position of the element in the periodic table satisfying the electronic configuration (n-1) d1 ns2 for n=4,

Ans.  (n-1) d1 ns2  = (4-1) d1 4s2

=  3d1 4s2

It lies in fourth period and III B group.


11. How does atomic size change in a group?

Ans.it increases from top bottom in a group.


12. Why Li and Mg show resemblance in chemical behaivour?

Ans. Due to diagonal relationship, since their atomic size, electro negativity and ionisation potential are almost the same.


13. The atomic radius of elements decreases along the period but Neon has highest size among III period element? Why

Ans. Ne is the only element in III period element which has Van der walls radius whereas the rest has covalent radius. And it is known fact that Van der walls radius is always greater than covalent radius.


14. Define valency.

Ans.  The combining capacity of an element is known as valency.


15. How does valency vary in a group and period in the periodic table?

Ans. In a group, the valency of an element remains constant while in a period it increases from left to right.


16. What is the valency of noble gases?

Ans. Noble gases on the extreme right are zero valent.


17. How do metals react in a period?

Ans.The tendency of an element to lose electrons decreases in going from left to right in a period. Thus the reactivity of metals goes on decreasing in a period from left right.


18. How do metals react in a group?
Ans.
 the tendency to lose electrons increases as we go down a group so the reactivity of metals increases down the group.


19. What is an amphoteric  oxide?
Ans.
 Oxides which behave as acids with bases and as a base with an acid are called amphoteric oxide.


20. Define a neutral oxide
Ans.
Neutral oxides have no acidic or basic properties.


21. What is the general outer electronic configuration of f – block elements?
Ans.
The general outer electronic configuration of f – block element is
(n – 2) f1-14 (n – 1) d0-11 ns2


22. Why do Na and K have similar properties?
Ans.
 Na and K have similar physical and chemical properties because they have same number of valence electrons.


23. Arrange the following elements in the increasing order of metallic character: Si, Be, Mg, Na, P.
Ans.
 P < Si < Be < Mg < Na


2 Marks Questions

1.How did Mendeleev arrange the elements?

Ans. Mendeleev arranged elements in horizontal rows and vertical columns of a table in order of their increasing atomic weights in such a way that the elements with similar properties occupied the same vertical column or group.


2.Name the two elements whose existence and properties were predicted by Mendeleev though they did not exist then.

Ans.   Mendeleev predicted not only the existence of gallium and germanium, but also described some of their general physical properties.


3.Give the main features of s-block elements.

Ans.S – block elements :- The elements in which the last electron enters the s – orbital of their outer most energy level are called s – block elements. It has elements of groups 1 and 2.  The general electronic configuration of s – block elements is ns1-2.


4.Give the main features of p-block elements.

Ans.p – block elements : The elements in which the last electron enters the p – orbital of their outermost energy level are called p – block elements. It contains elements of group 13,14, 15, 16, 17 and 18 of the periodic table. General electronic configuration of p – block elements is ns2 np1-6.


5.Give the main features of d-block elements.

Ans. d – block elements :- The elements in which the last electron enters the d – orbitals of their last but one energy level constitute d – block elements. There block consists of the elements lying between s and p block starting from 4th period and onwards. They constitute groups 3 to 12 in the periodic table. General electronic configuration is (n – 1) d1-10  ns1-2.


6.Give the main features of f-block elements.

Ans. f – block elements : The elements in which the last electron enters the f – orbital of their atoms are called f – block elements. In these elements the last electron is added to the third to the outermost energy level. These consist of two series of elements placed at the bottom of the periodic table known as Lanthanoid and actinoid series. General electronic configuration is (n-2)f1-14 (n-1)d0-1 ns2.


7.Explain why cations are smaller and anions are larger in radii than their parent atom?

Ans.The radius of cation is smaller than the parent atom. Cation is formed by the loss of one or more electron from the gaseous atom, but the nuclear charge remains the same. As a result, the nuclear hold on the remaining electrons increases because of the increases in the effective nuclear change per electron resulting in decrease in size.

Whereas anion is formed by the gain of one or more electrons by the gaseous atom but the nuclear charge is same though the number of electrons has increased. The effective nuclear charge per electron decrease in the anion and the cloud is held less tightly by the nucleus. This causes increase in size.


8.Define ionization enthalpy and electron gain enthalpy?

Ans.  Ionization enthalpy – It represents the energy required to remove an electron from an isolated gaseous atom (x) in ground state resulting in the formation of a positive ion.

x(g)  + Energy ® x+ (g) + e

Electron gain enthalpy – When an electron is added to a neutral gaseous atom (x) to convert it into a negative ion, the enthalpy change accompanying the process is defined as the electron gain enthalpy.

x(g) +  e  ® x (g)


9.How does metalic charecter change in a group?

Ans. It increases from top to bottom in a group


10.The size of an atom can be expressed by three radii. Name them. Which of these given the highest, and the lowest value of the atomic radius of an element?

Ans. The atomic size are generally expressed in terms of the following radii covalent radius, metallic radius and Van der waal’s radius.

Van der waal’s radius > Metallic radius > covalent radius.


11.Among the elements B, Al, C and Si 

(a)  Which has the highest first ionization enthalpy?

(b)  Which has the largest atomic radius?

Ans.

(a)   Carbon has the highest first ionization enthalpy.

(b)   Aluminum has the largest atomic radius.


12.Na+ has higher value of ionization enthalpy than Ne, though both have same electronic configuration.

Ans.  Na+ and Ne both has 10 electrons but Na+ having, 11 protons in its nucleus (Ne has 10 protons) exert higher effective nuclear charge and thus removal of electron from Na+ requires more energy.


13.How does the reactivity of non–metals changes in a period and group? 

Ans.The reactivity of non – metals is measured in terms of its tendency to gain electrons to form an ion. The reactivity of non – metals increases from left to right in a period whereas reactivity decreases in a group as we go down the group because the tendency to accept electrons decreases down the group.


14.Give the properties of the oxides in a particular period.

Ans.Elements on two extremes of a period easily combines with oxygen to form oxides. The normal oxide formed by the element on extreme left is the most basic (eg. Na2O) whereas that formed by the element on extreme right is the most acidic (eg. clO7). Oxides at the centre are however amphoteric  (eg. Al2 O3) or neutral (eg. CO).


15.Why does lithium form covalent bond unlike other alkali which forms ionic bond?     

Ans.Lithium forms covalent bond which is different from its group members because of its anomalous behaviour Li is small in size, large charge / radius ratio and has high electro negativity value. Also it has only 1s2 2s1 orbital for bonding.


16.The atomic number of an element is 16. Determine its position in accordance to its electronic configuration.

Ans. The atomic number of the element is 16.

The electronic configuration of the element is 1s2 2s2 2p6 3s2 3p4

Thus the element belongs to ‘p-block’ and is placed in third period and 16th group of the periodic table.


17. Why are elements at the extreme left and extreme right the most reactive?

Ans. The maximum chemical reactivity at the extreme left (among alkali metals) is exhibited due to the loss of an electron leading to the formation of a cation due to low ionization enthalpy and at the extreme right (among halogens) shown by the gain of an electron forming an anion. Due to high electron affinity.


18. Why does the ionization enthalpy gradually decreases in a group?

Ans. In a group, the increase in atomic and ionic radii with increase in atomic number generally results in a gradual decrease in ionization enthalpies.


19. Why does electronegativity value increases across a period and decreases down period?

Ans. The attraction between the outer electrons and the nucleus increases as the atomic radius decreases in a period. The electeonegativity also increases. On the same account electronegativity value decreases with the increase in atomic radii down a group.


20.How does electronegativity and non – metallic character related to each other?

Ans.Electronegativity is directly related to the non – metallic character of elements. Electronegativity is inversely related to the metallic properties of elements. Thus the increase in electronegativities across a period is accompanied by an increase in non – metallic properties of elements. Similarly, the decrease in electronegativity down a group is accompanies by a decrease in non – metallic properties of elements.

21.Describe the main features of Mendeleev’s periodic table?

Ans.(i)In Mendeleev table, the elements were arranged in vertical columns, and horizontal rows. The vertical columns were called groups and the horizontal rows were called periods.

(ii)There were in all eight groups. Group I to VIII. The group numbers were indicated by Roman numerals. Group VIII occupy three triads of the elements each i.e. in all nine elements.

(iii)There were seven periods to accommodate more elements the period 4, 5, 6 and 7 were divided into two halves. The first half of the elements were placed in the upper left corner and the second half in the lower right corner of each box.

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