Classification Of Elements And Periodicity In Properties

Radii of Mg²⁺< Na⁺ due to an increase in effective nuclear charge on the outermost electrons as we go down the group. Na⁺ and Mg²⁺ are cations so they have a smaller radii than F^- and O2^-.

In the case of F^- and O2^- fluoride has an extra proton as compared to oxygen and the number of electrons gained by oxygen is more which increases the nuclear charge on fluoride.

Terbium is a lanthanoid as it has z=65. Actinoids are the elements which have an atomic number in the range of 89-103.

As shown in the above periodic table Tb is a lanthanoid and is in that series.

when the electron is in the same shell then the order of the screening effect is s > p > d > f.

shielding effect is the decrease in effective nuclear charge on the electron cloud due to the inner shells which results in an easy removal of the valence shell electrons as nuclear attraction on them decreases.

Na configuration: 1s²,2s²,2p⁶,3s¹.

Mg configuration: 1s²,2s²,2p⁶,3s².

Al configuration: 1s²,2s²,2p⁶,3s²,3p¹

Si configuration: 1s²,2s²,2p⁶,3s²,3p²

As effective nuclear charge on Na is less than Mg, so the ionisation enthalpy of Na will be less than Mg as Mg has 2 electrons and Na has one electron in the outermost shell.

It is easier to remove an electron from p orbital than in s orbital so the ionisation enthalpy of Mg is less than Al. Therefore, Ionisation enthalpy of Mg is greater than Al.

As z= 64 so the nearest noble gases would be xenon. After that 10 electrons would be left so the configuration is given by this option. filling of s orbitals of nth shell takes before the filling of d orbitals of (n-1)th shell. Therefore, option ii) is incorrect. Options i) and iv) are incorrect as in the third option f orbital has half filled stability which results in the stability of the whole atom.

3d orbitals are filled after 4s in transition elements.

The modern periodic table is based on the properties that elements are a periodic function of their atomic numbers and once in a period an atom attains a stable configuration then the next element which occurs in the group shows similar properties as that of the element in the same group but of the previous period. This shows that option i) is correct.

The ionisation enthalpies increase as we move along the period due to increase in the effective nuclear charge. Thus option iv) is also correct.

This periodic table clearly shows that metals are more in number than non-metals. So option ii) is also correct.

Hence, the correct answer is (iii).

As we go down a group electron gain enthalpy decreases but in case of chlorine and fluorine, fluorine has a smaller size due to which it faces inter electronic repulsions. So, electron gain enthalpy of fluorine is less than that of chlorine.

The maximum number of shells are shown by the maximum principle quantum number which is equal to the period number of the periodic table.

For example, Electronic Configuration of Na is 1s²2s²2p²3s¹

As the last electron is present in the third shell, therefore sodium lies in the third period.

Lanthanoid are known for their 4 f orbitals so the electrons are also filled in these orbitals.

For example, [Xe] 4f⁷ 5d¹ 6s²= Electronic Configuration of Gadolinium

This is an element present iin the lanthanoid series and has electrons present in the 4f orbitals.

(Diagram of Lanthanoids).

I^- has a larger size as after gaining one electron its effective nuclear charge decreases. The size of anion is greater than the size of the neutral atom and the size of cation is always less than the neutral atom.

Ois an electronegative element so when we add an electron to it, it gains that electron and forms O^- anion. Thus energy is released in the formation and it is an exothermic process. But while another electron is added to this unstable O^- anion then it faces repulsion due to the similar charges. So in order to overcome these repulsions due to similar charge external energy is required due to which this becomes an endothermic process.

Electronic configuration of element with atomic number 57 is given by

[Xe]5d¹6s². This element is La.

The last electron of this element with atomic number 57 enters the d subshell so it is a d block element.

Rn is the last element of 6^th period and has an atomic number of 86. So the last electron enters in the 6p subshell. Thus option (iii) is correct.

Option i) and ii) are incorrect as the last electron enters the 7^th shell therefore they belong to 7^th period.

In case of option iv) the configuration shows that the element is not the last element as it would have a stable orbital.

The long form of periodic table can just accommodate 118 elements. So 126 cannot be accommodated in it.

Electronic configuration of element with atomic number 43 is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s³ 4p⁶4d⁵ 5s². As the last electron lies in the 5^th shell therefore this element belongs to the 5^th period. Element just above this belongs to the 4^th period with atomic number 43-18= 25.

From the periodic table and their electronic configuration is can be easily determined that they are halogens.

[Ar]4s² 4p⁵=35

[Kr]5s²5p⁵=53

[Xe]6s²6p⁵=85

As the p orbitals has 5 electrons and there are 7 electrons in the last shells therefore these are group 17 elements which are also called halogens.

The tendency to gain electrons depend upon the configuration and stability of an element. (A) is the most stable that is why least tendency and (D) needs one electron to become stable that is why highest electron gain enthalpy.

Vacant d-orbitals in elements allows transfer of electrons after unpairing electrons from s and p orbitals. D- orbitals are present in P and S due to which they can accommodate extra electrons which increases their covalency.

In excited state sulphur has a valency of six due to six unpaired electrons. Similarly in phosphorous it has a valency of 5 in the excited state due to the presence of the 5 unpaired electrons.

For example,

Sulphur configuration: 1s² 2s² 2p⁶ 3s² 3p⁴

Sulphur configuration excited state: 1s² 2s² 2p⁶3s¹3p³3d²

Phosphorous configuration: 1s² 2s² 2p⁶ 3s² 3p³

Phosphorous configuration excited state: 1s² 2s² 2p⁶3s¹ 3p³3d¹

Group 1 and 2 elements have low ionisation enthalpies as compared to elements of other group. When we provide energy to them then the electrons excite from a lower energy level to a higher. Once they reach the ground state then a particular wavelength is emitted imparting colour to them.

3 and 87 belongs to group 1, 33 belongs to group 15 and 53 belongs to group 17.

9 and 35 belongs to group 17, 51 to group 15 and 88 to group 2.

Chlorine gains electron more readily as it needs one electron to attain a stable configuration. Na is a metal so it donates electron more readily. As O has an inter-electronic repulsion so sulphur has a greater tendency to gain electron.

Helium has a small size and stable electronic configuration due to which it has the highest first ionization enthalpy in the periodic table.

The atomic radii of alkali is greater due to less effective nuclear charge on the outermost electrons due to their electron donating tendency.

i) Zn²⁺=30-2=28 Ca²⁺=20-2=18 Ga³⁺=31-3=28 Al³⁺=13-3=10

ii)K⁺= 19-1=18 Ca²⁺=20-2=18 Sc³⁺=21-3=18 Cl^-=22-4=18

iii)P^3- =15+3=18 S^2- =16+2=18 Cl^- =17+1=18 K⁺=19-1=18

iv)Ti⁴⁺= 22-4=18 Ar=18 Cr³⁺=24-3=21 V⁵⁺=23-5=18

the second and third are correct options as they have isoelectronic elements.

Electron gain enthalpy of fluorine is less than that of chlorine so third option is incorrect.

First ionisation of nitrogen is greater than that of oxygen. Thus ii) is incorrect.

Electron gain enthalpy and ionisation enthalpy have units of enthalpy but electronegativity and metallic character don’t have.

As effective nuclear charge increases ionic radii decreases as it causes a huge attraction the outermost shell.

Screening effect shield the outermost electron and is directly proportional to ionic radii.

An element that belongs to 3^rd period and group 13 is aluminium. It is a solid and non metallic element. It is non metallic due to high ionisation energy.

Fluorine has a smaller size as compared to chlorine as a result of which the attraction outside the shell to gain electron is less. Moreover there are inter electronic repulsions as well in the 2p orbitals which results in the less negative electron gain enthalpy.

Those elements which have their outermost electrons in the d orbitals are known as d block elements. But to be a transition element it is important to have an incompletely filled d orbital of the element or their respective ions. For example, zinc, cadmium and mercury.

There are 118 elements in the 7 periods of the modern periodic table. Therefore the element with atomic number 119 will lie in the 8^th period of the first group and will have the outermost electronic configuration of 8s¹. It belongs to group 1 and has a valency of one. The formula of its oxide will be M₂O.

The ionisation enthalpy of elements varies across period and group. The ionisation enthalpy decreases down a group and increases as we move from left to right in a period. It is affected by the following parameters.

• Effective nuclear charge on the outermost electrons.

• Electron- electron repulsion force.

• Stability of the element due to half filled and fully filled orbitals.

But this trend is not true for every element for example beryllium and boron. Due to stable configuration of Be a very high first ionisation energy is required to remove an electron but in case of boron it has one electron in 2p¹ less energy is required.

Electronic configuration of Nitrogen is 1s² 2s²2p_x¹2p_y¹2p_z¹

Electronic configuration of Oxygen is 1s² 2s²2p_x²2p_y¹2p_z¹

As nitrogen has half filled stable p orbital therefore it has a positive electron gain enthalpy.

The ionisation enthalpy of oxygen is lower than that of Nitrogen as because when we remove one electron from oxygen then it easily donates it to attain half filled stability but in case of nitrogen it is difficult to remove one electron as it already has half filled stability and it would become unstable after that.

(i) Carbon has the highest ionisation enthalpy as The ionisation enthalpy of elements varies across period and group. The ionisation enthalpy decreases down a group and increases as we move from left to right in a period. It is affected by the following parameters.

(ii) Aluminium has the most metallic character.

Metallic character increases on moving down the group and decreases on moving left to right in a period.

1. They show variable oxidation states. The reducing character increases down the group and oxidising character increases along the period.

2. They have a high ionisation enthalpy than the s-block elements.

3. They usually form covalent compounds.

4. Both metals and non metals are found in this group, but non metals are slightly more in number. As we go along the period non metal character increases and down the group metallic character increases.

(i) 72, 160

EXPLANATION

Neon has vanderwaals radii and fluorine has covalent radii. Covalent radius is always less than vanderwaal’s radius, so the radius of Fluorine is 72pm and Neon is 160pm.

Ti has an atomic number off 22 and electronic configuration [Ar]3d²4s² and can show three oxidation states of +2,+3 and +4 in various compounds like TiO₂9(+4), Ti₂O₃(+3) and TiO(+2). The non transition elements like the p block elements show variable oxidation states like in case of phosphorous. It has -3,+3 and +5. Lower oxidation states are ionic in nature as they accept electrons and higher oxidation states by unpairing the electrons and shifting to d orbitals.

Electronic configuration of Nitrogen is 1s² 2s²2p_x¹2p_y¹2p_z¹

Electronic configuration of Oxygen is 1s² 2s²2p_x²2p_y¹2p_z¹

As nitrogen has half-filled stable p orbital therefore it has a positive electron gain enthalpy.

The ionisation enthalpy of oxygen is lower than that of Nitrogen as because when we remove one electron from oxygen then it easily donates it to attain half-filled stability but in case of nitrogen it is difficult to remove one electron as it already has half-filled stability and it would become unstable after that.

First member of each group of representative elements (i.e., s and p-block elements) shows anomalous behaviour due to the following reasons:

• Due to small atomic size

• D-orbitals are absent.

• Ionisation enthalpy

• Electron gain enthalpy

Li is a first group element. It has different properties and forms covalent compounds and nitrides.

Beryllium is first element of second group. It has various anomalies like it forms a covalent compound with coordination number four unlike other elements that have a coordination number 6.

ACIDIC OXIDES

SO₂ , B₂O₃ are acidic oxides and p block elements.

Reaction of SO₂ with water

SO₂ + H₂O → H₂SO₃

Reaction of B₂O₃ with water

B₂O₃ +3 H₂O → 2H₃BO₃

Acidic oxides are those oxides that form acids after reacting with water.

BASIC OXIDES

CaO, BaO, Ti₂O form basic oxides

Reaction of Ti₂O with water.

BaO + H₂O → Ba(OH)₂

REACTION OF CaO WITH WATER

CaO + H₂O → Ca(OH)₂

Basic oxides are those oxides that form bases after reacting with water

AMPHOTERIC OXIDES

Zinc oxides and aluminium oxides are the two amphoteric oxides.

REACTION OF ZnO WITH WATER

ZnO + 2H₂O + 2NaOH → Na₃Zn[OH]₄ + H₂

ZnO +2HCl → ZnCl₂ + H₂O

REACTION OF Al₂O₃ WITH WATER

Al₂O₃ (s) + 6 NaOH(aq) + 3H₂O(l) → Na₃ [Al(OH)₆] (aq)

Al₂O₃ (s) + 6HCl(aq) + 9H₂O(l) → 2[Al(H₂O)₆]³⁺(aq) + 6Cl^-

Those oxides that form both acids and bases are known as Amphoteric Oxides.

Electronic configuration of Na=1s² 2s²2p⁶3s¹

Electronic configuration of Mg=1s² 2s²2p⁶3s²

After losing one electron sodium attains stable configuration that is why its first ionisation enthalpy is less than that of magnesium. But in case of second ionisation enthalpy magnesium has one electron in its outermost shell and to attain stability it loses that easily as compared to sodium which is already stable.

Exothermic reaction- The reaction in which heat is evolved is called exothermic reaction.

for example,

CaO + CO₂→ CaCO₃ ΔH=-178kJmol^-1

Endothermic reaction- The reaction in which heat is absorbed is called endothermic reaction.

for example,

2NH₃ → 3N₂+ H₂ ΔH=918kJmol^-1

(i) increasing first ionisation enthalpy.

EXPLANATION

S< P< O< N is the increasing order of the first ionisation enthalpy.

As we go down the group the ionisation enthalpy decreases and as we go along the period then it increases but in case of oxygen and nitrogen due to half-filled stability of 2 p orbitals of nitrogen it has a higher ionisation enthalpy than oxygen.

(ii) increasing non-metallic character.

P<S<N<O

EXPLANATION

As we go down the group non-metallic character decreases as effective nuclear charge on the outermost shell decreases which helps to gain an electron. As we move along the period then effective nuclear charge increases which increases the non-metallic character.

The ionisation enthalpy of elements varies across period and group. The ionisation enthalpy decreases down a group and increases as we move from left to right in a period. It is affected by the following parameters.

• Effective nuclear charge on the outermost electrons.

• Electron- electron repulsion force.

• Stability of the element due to half-filled and fully filled orbitals.

As we move from left to right in a period, the size of the atoms decreases due to the increase in the effective nuclear charges on the outermost electron. As a result of which electronegativity of elements increases on moving from left to right in the periodic table.

(b) As we go down the group the atomic size increases which results in the increase in the distance of the electrons in the outermost shell as a result of which effective nuclear charge decreases. This results in the decrease of ionisation enthalpy.

When we move from left to right in a period then the metallic character decreases and non-metallic character increases as there is an increase in ionisation enthalpy and electron gain enthalpy along the period.

Sodium atom loses one electron to form sodium cation and after the formation of a cation the effective nuclear charge on the ion increases on the left electrons which results a decrease in radius.

Caesium is the least electronegative alkali metal as electronegativity decreases as we go down the group. Caesium is a group 1 element and lies down the group.

Electronegativity increases as we move along the period and it decreases as we move down the group. So this states that group one elements are least electronegative elements in a period and the element at the bottom of group 1 will be the least in the group. Thus caesium has the highest electronegativity as it has the largest size due to decrease in the effective nuclear charge.

EXPLANATION

When we move from left to right in a period, effective nuclear charge of elements increases which decreases the atomic radius.

(i) Most reactive non metal is B

REASON: It has the highest negative electron gain enthalpy.

(ii) Most reactive metal =A

REASON: It has the lowest negative electron gain enthalpy as well as lowest first ionisation energy.

(iii) Least reactive element=D

REASON:D is the least reactive element as it has an extremely high first and second ionisation enthalpy.

(iv) Metal forming binary halide=C

REASON: it has low first ionisation enthalpy than second but the different is not much.

EXPLANATION

Electron gain enthalpy of elements having fully filled orbitals will be extremely positive as they do not want to lose their stability. So in the first case it is +48kJmol^-1. In the second case s orbital has just one electron so this element has more tendency to donate electron but as it can also gain one electron to complete its s orbital so it has a low negative electron gain enthalpy. In the third case just one electron is required to fill the p orbital. Hence it has a very high negative electron gain enthalpy. In the fourth case as 2 more electrons are required so it has a negative electron gain enthalpy but it is less negative than the third case.

(ii) Assertion and reason both are correct statements and reason is correct explanation of assertion.

EXPLANATION

Ionisation enthalpy increases from left to right in a period due to increase in effective nuclear charge along a period.

(i) Assertion and reason both are correct statements and reason is correct explanation for assertion.

Boron has [He]2s²2p¹ configuration and after donating one electron it attains stable configuration therefore it has smaller first ionisation enthalpy than beryllium which has a completely filled s orbital as it a configuration [He]2s². As 2p electrons are shielded by the inner electrons so it is easy to remove an electron from these orbitals and their ionisation energy decreases.

(iv) Assertion is wrong statement but reason is correct statement.

Electron gain enthalpy becomes less negative as we go down the group as on going down the atomic size increases due to the increase in the number of shells which makes electron farther.

The factors affecting electron gain enthalpy and the trend in its variation in the periodic table are:

1)ATOMIC SIZE

As we go down the group electron gain enthalpy decreases as the distance of the nucleus from the outermost shell increases which decreases its tendency to gain electron and electron gain enthalpy becomes less negative.

2)EFFECTIVE NUCLEAR CHARGE

As we go from left to right in a period the effective nuclear charge increases and when we move down the group it decreases which results in the attraction of electrons from the outermost shell.

3) ELECTRONIC CONFIGURATION

The tendency to gain electrons depends upon the stability of an element. Elements having completely or half filled stable orbitals have very low tendency to gain electron thus they have very low electron gain enthalpy.

TRENDS

Across a period: electron gain enthalpy becomes more negative.

Down the group: electron gain enthalpy becomes less negative.

Ionisation enthalpy is the energy required by an isolated and gaseous atom in its ground state to remove an electron.

A(g) A⁺(g) →+e^-

Factors affecting ionisation enthalpy

Effective nuclear charge

Due to the screening effect the valence electrons are shielded by the inner core electrons. This effective nuclear charge is less than the actual charge present on the atom.

Penetrated orbitals

It is difficult to remove an electron from the orbitals that are closer to the nucleus and are penetrated towards the nucleus. The order of penetration is given by:

s>p>d>f

stability of orbitals

Half filled and fully filled orbitals have a high ionisation enthalpy as they don’t want to lose their stability.

Across a period: ionisation enthalpy increases along the period.

Down the group: Ionisation enthalpy decreases down the group.

In the outermost shells the electrons are distributed as per the atomic numbers due to which similar properties occur. As we go down the group or along the period there are many chemical properties which depend on the atomic number.

For example,

Electronic configuration of group 1 elements is:

1s²2s¹ =Li₃

1s²2s²2p⁶3s¹ = Na₁₁

1s²2s²2p⁶3s²3p⁶4s¹=K₁₉

1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 5s¹=Rb₃₇

1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s²5p⁶6s¹=Cs₅₅

[Rn] 7s¹=Fr₈₇

Number of electrons in the valence shell determine the group of an element. All these elements have one electron in their valence shell.

Electronic configuration of group 1 elements is:

1s²2s¹ =Li

1s²2s²2p²3s¹ = Na

1s²2s²2p²3s²3p⁶4s¹=K

1s² 2s² 2p² 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 5s¹=Rb

1s² 2s² 2p² 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s²5p⁶6s¹=Cs

[Rn] 7s¹=Fr

Number of electrons in the valence shell determine the group of an element. All these elements have one electron in their valence shell.

• Some elements do not fit to his classification that was based on the atomic weights. For example, cobalt was placed before nickel, tellurium before iodine etc.

• He violated his periodic law at certain places to assign the correct position to elements on the basis of their properties.

• The position of hydrogen was not defined correctly as it resembled few characteristics of alkali metals and few of halogens.

• Position of isotopes would be different according to Mendeleev’s periodic table.

• No separate positions were given to lanthanoids and actinoids.

• Position of group VIII elements was not justified.

• The properties be it physical or chemical of an elements are periodic functions of their atomic numbers is the modern periodic law. According to Mendeleev’s these properties were a periodic function of their atomic masses. Atomic number is a better parameter which makes modern periodic table superior. For example Iodine having lower atomic mass than tellurium was placed ahead of Tellurium in the mendeleev’s periodic table because Iodine showed similar properties with other halogens.

• Metals, non metals and metalloids can be easily differentiated in the modern periodic table.for example the non metals are present on the upper right side of the table.

• Position and arrangement of elements can be easily determined by their electronic configurations. For example elements that have one electron in their last shell are placed in group 1.

• Each vertical column that is a group posses similar properties.

• The modern periodic table allows the systematic filling of electrons along a period. There is an increase of one electron until a stable noble gas configuration is reached.

• The modern periodic table shows increase in energy shells down the group.

• The 8^th group has been allotted a correct position.

• Transition elements are placed between s and p block elements as the properties shown by them lie in between the two of them.

• Unknown elements can be easily located according to their atomic numbers. For example in mendeleev’s periodic table he has to left some gaps to place gallium and germanium.

The ionisation enthalpy decreases on moving down a group and increases as we move along a period

1)ATOMIC SIZE

As we go down the group effective nuclear charge decreases as the distance of the nucleus from the outermost shell increases which decreases its tendency to lose electron. As we go along the period the ionisation enthalpy increases as effective nuclear charge increases. The increase in number of shells result in decrease in ionisation enthalpy.

2)EFFECTIVE NUCLEAR CHARGE

As we go from left to right in a period the effective nuclear charge increases and when we move down the group it decreases which results in the attraction of electrons from the outermost shell.

3) ELECTRONIC CONFIGURATION

The tendency to gain electrons depends upon the stability of an element. Elements having completely or half filled stable orbitals have very low tendency to gain electron thus they have very low electron gain enthalpy.

Group 1 elements have the lowest ionisation enthalpy as they have just one electron in the outermost shell and as we go down the group this ionisation enthalpy decreases due to increase in atomic size. As we go along the period from group 1 to group 17 elements then ionisation energy increases as they have 7 electrons in the outermost shell and a high effective nuclear charge on the outermost shell due to which it is difficult to lose an electron.