Preview:  

We shall begin the chapter by discussing the various methods used by scientists like, Dobereiner, de Chancourtois, Alexander and Mendeleev to classify elements in certain order such that they exhibit periodicity when placed in that specific order. Periodicity is the recurrence or repetition in properties of elements with increasing atomic number.

Then, we shall see why Mendeleev’s periodic table was the most acceptable form of classification of elements.

Next, we shall see how Mendeleev’s periodic table was modified into The Modern Periodic Table, which explained more accurately, the cause of periodicity of elements.

In the subsequent sections, we shall study certain trends exhibited by elements in the periodic table, which includes: electronic configuration of elements, metallic property, atomic radius, ionic radius, ionization enthalpy, electron affinity and valency.

Lastly, we shall study the anomalous behaviour of elements belonging to certain groups.

We shall come to the end of the chapter with the study of reducing and oxidizing nature of metals and non-metals.

3.0 Introduction:

In this chapter, we shall deal with Periodic Table and the Modern Periodic Law.

We will also learn how the periodic classification follows as a logical consequence of the electronic configuration of atoms. Some of the periodic trends in the physical and chemical properties of the elements shall also be studied.

3.1 Why do we need to classify elements?

• With large number of elements, it is very difficult to study individually the chemistry of all these elements and the compounds they form. To ease out this problem, scientists systematically began to organise their knowledge by classifying the elements.
• Classification is also essential to justify known chemical facts about elements and even predict new ones for further study.

3.2 Genesis of Periodic Classification

Dobereiner’s law of triads:

• The German chemist, Johann Dobereiner started classification of elements based on similarity in physical and chemical properties of elements in groups of three (triad).

In each of these triads, he noticed that:

• the middle element of each of them had an atomic weight about half way between the atomic weights of the other two (Table 3.1).

• the properties of the middle element were in between those of the other two members.

• This pattern was called Law of Triads and it seemed to work only for a few elements. Hence, it was dismissed as coincidence.

de Chancourtois’s classification:

French geologist, A.E.B. de Chancourtois arranged the then known elements in order of increasing atomic weights and made a cylindrical table of elements to display the periodic recurrence of properties. This also did not attract much attention.

Alexander’s law of octaves:

• The English chemist, John Alexander Newlands propounded the Law of Octaves.
• He arranged the elements in increasing order of their atomic weights and noted that every eighth element had properties similar to the first element (Table 3.2). The relationship was just like the octaves of music in which every eighth note resembles the first.

• Limitation: Newlands’s Law of Octaves seemed to be true only for elements up to calcium.

Note box: Although his idea was not widely accepted at that time, he, for his work, was later awarded Davy Medal in 1887 by the Royal Society, London.

Introduction to Mendeleev and Lothar Meyer’s classification:

The Russian chemist, Dmitri Mendeleev and the German chemist, Lothar Meyer working independently, both proposed that, “on arranging elements in the increasing order of their atomic weights, similarities appear in physical and chemical properties at regular intervals”.

Lothar Meyer plotted the physical properties such as atomic volume, melting point and boiling point against atomic weight and obtained a periodically repeated pattern. Although he had developed a periodic table similar to that of the modern periodic table, his work wasn’t published until after the work of Dmitri Mendeleev, the scientist who is generally credited for the development of the Modern Periodic Table.

Concept box: Periodicity refers to trends or recurrence or repetition in properties of elements with increasing atomic number.

Definition box: Mendeleev’s periodic law: It states that, the properties of the elements are a periodic function of their atomic weights.

Mendeleev’s periodic table:

• Mendeleev arranged elements in horizontal rows and vertical columns in 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.
• He gave more importance to the similarities in properties of elements which were periodic functions of their atomic weights. In particular, Mendeleev relied on the similarities in the empirical formulas and properties of the compounds formed by the elements.
• Mendeleev’s periodic table consists of:

a) Nine vertical columns known as groups –

Groups I to VII were subdivided into subgroups: “A” and “B”;

Group VIII consists of triads and;

Group IX is called “0-group” and is comprised of insert gases.

b)   Seven Horizontal rows, known as periods.

• He realized that, some of the elements did not fit in, if the order of atomic weight was strictly followed. Hence, he ignored the order of atomic weights, thinking that the atomic measurements might be incorrect, and placed the elements with similar properties together. For example, iodine with lower atomic weight than that of tellurium (Group VI) was placed in Group VII along with fluorine, chlorine, bromine because of similarities in properties (refer Figure 3.1 above).
• Keeping his primary aim of having the elements with similar properties in the same group, he proposed that some of the elements were still undiscovered and, therefore, left several gaps in the table. For example, both gallium and germanium were unknown at the time Mendeleev published his Periodic Table. He left the gap under aluminium and a gap under silicon, and called these elements Eka- Aluminium and Eka-Silicon.
• Mendeleev not only predicted the existence of gallium and germanium, but also described some of their general physical properties. These elements were discovered later.

Some of the properties predicted by Mendeleev for these elements and those found later, experimentally are listed in Table 3.3.

Merits of Mendeleev’s periodic table:

• It helped in classification of elements with similar properties into one group.
• Helped in discovery of new elements because Mendeleev predicted the existence of these elements. Mendeleev also predicted the properties of elements on the basis of their positions.
• It is useful in correcting doubtful atomic weights of certain elements, predicting valency of elements and in assigning correct position to certain elements.

Defects in Mendeleev’s periodic table:

• Position of Hydrogen: Hydrogen shows similarity in certain chemical properties with halogens and alkali metals and Mendeleev could not decide whether it should be placed at the top of group I (Alkali metals) or on the top on group VII (Halogens).
• Position of Isotopes: Isotopes of an element have similar chemical properties but differ in atomic masses. The isotopes thus occupy same group. But on the basis of atomic masses, they should occupy different groups. Hence, Mendeleev could not explain the positions of isotopes.
• Cause of Periodicity: Mendeleev could not explain why elements exhibit a periodicity in their properties when arranged in order of increasing atomic masses.
• No place was assigned to noble gases.
• Some elements with dissimilar properties were placed under same group and those with similar properties were placed under different groups.

Story box: Dmitri Mendeleev taught at the University of St.Petersburg where he was appointed Professor of General Chemistry in 1867. Preliminary work for his great textbook “Principles of Chemistry” led Mendeleev to propose the Periodic Law and to construct his Periodic Table of elements. The discovery of the first two noble gases helium and argon in 1890 suggested the possibility that there must be other similar elements to fill an entire family. This idea led Ramsay to his successful search for krypton and xenon. Work on the radioactive decay series for uranium and thorium in the early years of twentieth century was also guided by the Periodic Table. Mendeleev was a versatile genius. He worked on many problems connected with Russia’s natural resources. He invented an accurate barometer. In 1890, he resigned from the Professorship. He was appointed as the Director of the Bureau of Weights and Measures. He continued to carry out important research work in many areas until his death in 1907. Mendeleev’s name has been immortalized by naming the element with atomic number 101, as Mendelevium. This name was proposed by American scientist Glenn T. Seaborg, the discoverer of this element, “in recognition of the pioneering role of the great Russian Chemist who was the first to use the periodic system of elements to predict the chemical properties of undiscovered elements, a principle which has been the key to the discovery of nearly all the trans-uranium elements”.

Questions from sections: 3.1 and 3.2:

1. Describe:

a) Dobereiner’s law of triads

b) de Chancourtois’s classification

c) Alexander’s law of octaves

2. What is periodicity?

3. State Mendeleev’s periodic law and explain the features of Mendeleev’s Periodic Table.

4. List the merits and defects in Mendeleev’s periodic table.

3.3 Modern Periodic Law and the Present Form of the Periodic Table

The English physicist, Henry Moseley observed regularities in the characteristic X-ray spectra of the elements by plotting “ν” (frequency of X-rays emitted) against atomic number, “Z” which gave a straight line and; the plot of “ν” vs atomic mass did not. He thereby showed that, the atomic number is a more fundamental property of an element than its atomic mass.

Mendeleev’s Periodic Law was, therefore was accordingly modified and is now known as, the Modern Periodic Law which can be stated as:

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

Cause of periodicity:

The recurrence of outer electronic configurations at regular intervals, determines the physical and chemical properties of elements and their compounds according to the law.

Modern Periodic Table:

Figure 3.2 Long form of the Periodic Table of the Elements with their atomic numbers and ground state outer electronic configurations. The groups are numbered 1-18 in accordance with the 1984 IUPAC recommendations. This notation replaces the old numbering scheme of IA–VIIA, VIII, IB–VIIB and 0 for the elements.

• The modern version, also called “long form” of the Periodic Table of the elements (Figure 3.2), is the most convenient and widely used.
• The horizontal rows (which Mendeleev called series) are called periods and the vertical columns, groups.
• Arrangement of Groups: Elements having similar outer electronic configurations in their atoms are arranged in same groups or families. According to International Union of Pure and Applied Chemistry (IUPAC), the groups are numbered from 1 to 18 replacing the older notation of groups IA … VIIA, VIII, IB … VIIB and 0.
• Arrangement of Periods: There are altogether seven periods. The period number corresponds to the highest principal quantum number (n) of the elements in the period. The first period contains 2 elements. The subsequent period consists of 8, 8, 18, 18 and 32 elements, respectively. The seventh period is incomplete and would have a theoretical maximum (on the basis of quantum numbers) of 32 elements.
• Lanthanoids and Actinoids: 14 elements of both sixth and seventh periods (lanthanoids and actinoids, respectively) are placed in separate panels at the bottom.

Questions from section 3.3:

1. State the modern periodic law and the cause of periodicity according to the law.

2. Explain the features of Modern periodic table.

3.4 Nomenclature of Elements with Atomic Numbers > 100

IUPAC nomenclature was introduced to name elements with atomic numbers greater than 100 according to the following rules:

• Names are derived directly from the atomic number of the element using the numerical roots for 0 and numbers, 1-9 by adding suffix: “ium”. (refer table 3.4 below)

• In certain cases, names are shortened. For example, bi-ium and tri-ium are shortened to bium and trium and en-nil to ennil.
• The symbol of the element is then obtained from the first letter of the roots of numbers which makes up the atomic number of the element. The IUPAC names for elements with atomic number (Z) above 100 are shown in Table 3.5 below.

• Thus, the new element first gets a temporary name, with symbol consisting of three letters. Later permanent name and symbol are given by a vote of IUPAC representatives from each country. The permanent name might reflect the country (or state of the country) in which the element was discovered, or pay tribute to a notable scientist.

Note box: As of now, elements with atomic numbers up to 118 have been discovered.

Example 3.1: What would be the IUPAC name and symbol for the element with atomic number 120?

Solution:

From Table 3.4, the roots for 1, 2 and 0 are un, bi and

3.5 Electronic Configurations of Elements and the Periodic Table

The outermost shell of an atom is known as valence shell. Based on the valence shell electronic configuration and the type of orbital in which the differential (last) electron enters, the elements in periodic table are divided into four blocks of elements, as: s, p, d and f block elements.

(a) Electronic Configurations in Periods

• Successive period in the Periodic Table is associated with the filling of the next higher energy shells or successive principal quantum numbers.
• The number of elements in each period is twice the number of atomic orbitals available in the energy level that is being filled.
• The first period (n = 1), starts with the filling of the lowest level orbital (1s) and therefore has two elements — hydrogen ($1s^1$) and helium ($1s^2$). This marks the completion of first shell (K).
• The second period (n = 2) starts with lithium in which, the third electron enters the 2s orbital. The next element, beryllium has four electrons and has the electronic configuration $1s^22s^2$. The next element boron fills the 2p orbitals with electrons. L shell is completed at neon ($2s^22p^6$). Thus there are 8 elements in the second period.
• The third period (n = 3) begins at sodium, and the added electron enters a 3s orbital. Successive filling of 3s and 3p orbitals gives rise to the third period of 8 elements from sodium to argon.
• The fourth period (n = 4) starts at potassium, and the added electrons fill up the 4s orbital. But now, before the 4p orbital is filled, filling up of 3d orbitals becomes energetically favourable.  These elements constitute for the 3d transition series of elements (that is, from scandium to zinc)—Scandium (Z = 21) with an electronic configuration: $3d^14s^2$ until zinc with an electronic configuration of: $3d^{10}4s^2$, thus filling the 3d orbitals completely. The period ends at krypton with the filling up of the 4p orbitals. Altogether there are 18 elements in this fourth period.
• The fifth period (n = 5) beginning with rubidium is similar to the fourth period and contains the 4d transition series comprised of yttrium (Z = 39) up until cadmium (Z = 48). This period ends at xenon with the filling up of the 5p orbitals.
• The sixth period (n = 6) contains 32 elements and successive electrons enter 6s, 4f, 5d and 6p orbitals. The 6s orbital starts getting filled from caesium (Z = 55) and is completely filled by lanthanum (Z = 57). The 4f orbitals begin to fill with cerium (Z = 58) and ends at lutetium (Z = 71) to give the 4f-inner transition series which is called the lanthanoid series.
• The seventh period (n = 7) is similar to the sixth period with the successive filling up of the 7s, 5f, 6d and 7p orbitals and includes most of the man-made radioactive elements. This period ends at the element with atomic number 118 which would belong to the noble gas family. Filling up of the 5f orbitals after actinium (Z = 89) up until lawrencium (Z = 103) gives the 5f-inner transition series known as the actinoid series.
• The 4f and 5f-inner transition series of elements are placed separately in the Periodic Table to maintain its structure and to preserve the principle of classification by keeping elements with similar properties in a single column.

Example 3.2: How would you justify the presence of 18 elements in the 5th period of the Periodic Table?

Solution: When n = 5, l = 0, 1, 2, 3. The order in which the energy of the available orbitals 4d, 5s and 5p increases is 5s < 4d < 5p. The total number of orbitals available are 9. The maximum number of electrons that can be accommodated is 18; and therefore 18 elements are there in the 5th period.

(b) Group-wise Electronic Configurations

Elements in the same vertical column or grouphave similar valence shell electronicconfigurations, the same number of electrons in the outer orbitals, and similar properties.

For example, the Group 1 elements (alkali metals) all have $ns^1$ valence shell electronic configuration as shown below:

Questions from sections, 3.4 and 3.5:

1. What is a valence shell?

2. Explain the trend in electronic configuration for elements across the period.

3.6 Electronic Configurations and Types of Elements s-, p-, d-, f- Blocks:

s-block elements:

• The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) are known as s-Block Elements.
• General valence shell electronic configuration of group 1 and group 2 elements is $ns^1$ and $ns^2$ respectively where “n” is principal quantum number of outermost shell.
• They are all reactive metals.
• They have low ionization enthalpies.
• They lose the outermost electron(s) readily to form +1 (in the case of alkali metals) or +2 (in the case of alkaline earth metals) oxidation states.
• The metallic character and the reactivity increase as we go down the group.
• Because of high reactivity, they are never found pure in nature.
• The compounds of the s-block elements, with the exception of those of lithium and beryllium are predominantly ionic.

3.6.2 The p-Block Elements

• The elements in which, the outermost electron enters p-orbital are known as p-block elements.
• The p-Block Elements comprise of groups: 13 to 18.
• The p-Block elements together with the s-Block elements are called Representative Elements or Main Group Elements.
• The outermost electronic configuration for p-Block elements varies from $ns^2np^1$ to $ns^2np^6$ across the period.
• All the orbitals in the valence shell of the noble gases are completely filled by electrons and it is very difficult to alter this stable arrangement by the addition or removal of electrons. The noble gases thus exhibit very low chemical reactivity.
• Preceding the noble gas family are two chemically important groups of non-metals. They are the halogens (Group 17) and the chalcogens (Group 16). These two groups of elements have highly negative electron gain enthalpies and readily add one and two electrons respectively to attain the stable noble gas configuration.
• The non-metallic character increases as we move from left to right across a period and metallic character increases as we go down the group.

3.6.3 The d-Block Elements (Transition Elements)

• The elements from Group 3 to 12 in the centre of the Periodic Table belong to the d-block. They are also known as transition elements because they form a bridge between the chemically active metals of s-block elements and the less active elements of Groups 13 and 14.
• These are characterised by the filling of inner d-orbitals by electrons and are therefore referred to as d-Block Elements.
• These elements have the general outer electronic configuration: $(n-1) d^{1-10}ns^{0-2}$.
• They are all metals.
• They mostly form coloured ions.
• They exhibit variable valence (oxidation states), paramagnetism and are thus, often used as catalysts.
• However, Zn, Cd and Hg which have the electronic configuration,$ (n-1) d^{10}ns^2$ do not show most of the properties of transition elements.

3.6.4 The f-Block Elements

• They are also called, Inner-Transition Elements.
• The elements are placed in two rows at the bottom of the Periodic Table, called the Lanthanoids, Ce(Z = 58) – Lu(Z = 71) and Actinoids, Th(Z = 90) – Lr (Z = 103).
• They are characterised by the outer electronic configuration: $(n-2) f^{1 – 14}(n-1) d^{0–1 }ns^2$.
• The last electron added to each element is filled in f-orbital.
• They are all metals.
• Actinoid elements exhibit large number of oxidation states.
• Actinoid elements are also radioactive.
• The elements after uranium are called Transuranium Elements.

Example 3.3: The elements Z = 117 and 120 have not yet been discovered. In which family / group would you place these elements and also give the electronic configuration in each case.

Solution: We see from Fig. 3.2, that element with Z = 117, would belong to the halogen family (Group 17) and the electronic configuration would be [Rn] $5f^{14}6d^{10}7s^27p^5$. The element with Z = 120, will be placed in Group 2 (alkaline earth metals), and will have the electronic configuration [Uuo]8$s^2$.

3.6.5 Metals, Non-metals and Metalloids

The elements can also be divided into Metals and Non-Metals.

Metals:

• Metals comprise more than 78% of all known elements and appear on the left side of the Periodic Table.
• Metals are usually solids at room temperature [mercury is an exception].
• Metals usually have high melting and boiling points [gallium and caesium have very low melting points (303K and 302K, respectively)].
• Metals are good conductors of heat and electricity.
• They are malleable (can be flattened into thin sheets by hammering) and ductile (can be drawn into wires).

Non-metals:

• Non-metals are located at the top right hand side of the Periodic Table.
• Non-metals are usually solids or gases at room temperature with low melting and boiling points (boron and carbon are exceptions).
• Non-metals are poor conductors of heat and electricity.
• Most non-metallic solids are brittle and are neither malleable nor, ductile.
• The elements become more metallic as we go down a group; the non-metallic character increases as one goes from left to right across the Periodic Table.

Metalloids:

The change from metallic to non-metallic character is not abrupt as shown by the thick zig-zag line in Figure 3.3. For example, elements like: silicon, germanium, arsenic, antimony and tellurium bordering this line and running diagonally across the Periodic Table show properties that are characteristic of both metals and non-metals. These elements are called Semi-metals or Metalloids.

Example 3.4: Considering the atomic number and position in the periodic table, arrange the following elements in the increasing order of metallic character : Si, Be, Mg, Na, P.

Solution: Metallic character increases down a group and decreases along a period as we move from left to right. Hence the order of increasing metallic character is: P < Si < Be < Mg < Na.

Questions from section 3.6:

1. Explain the trends in electronic configuration group-wise (for, s, p, d and f blocks).

2. Explain briefly, the characteristics of:

a) Metals

b) Non-metals

c) Metalloids

3.7 Periodic Trends in Properties of Elements

Elements of the $18^{th}$ group which have completely filled outermost energy level ($ns^2np^6$) are stable. Rest of the elements tend to attain the configuration of their nearest 18^th group element, by losing or by gaining or sharing of electrons present in the outermost energy level. Elements with few electrons in the outermost shell (1^st and 2^nd group) tend to lose electrons and hence, they are said to be Electropositive. These elements are metals.

Those with larger number of electrons (belonging to 16^th and 17^th group) in the outermost shell have a tendency to gain electrons and hence, they are said to be, Electronegative. Generally, they are non-metals.

Rest of the elements tend to share electrons. Most of these are metals and few of them are non-metals.

However, the exact behaviour of these elements depends upon the atomic radius, ionization energy, electron affinity and electronegativity. These properties are called, Periodic properties.

3.7.1 Trends in Physical Properties

(a) Atomic Radius

• For a homo-nuclear diatomic molecule, atomic radius is equal to “half of the inter-nuclear distance between two identical atoms bonded by single covalent bond”.
• Atomic radius = AB/2

• It is expressed in nm or pm.
• For a given atom, the atomic radius is, “the distance between the centre of nucleus and the point at which the electron density is negligible”.

Variation of atomic radius in the periodic table:

• The atomic radii of a few elements are listed in Table 3.6.

• Across a period, the added electron enters the orbitals of the same principal shell but, the effective nuclear charge increases. The increase in effective nuclear charge increases the force of attraction between the electrons and the nucleus. Hence, the atomic radius decreases across a period.

Figure 3.4 (a) Variation of atomic radius with atomic number across the second period

• For alkali metals and halogens, as we descend the groups, the principal quantum number (n) increases and the valence electrons are farther from the nucleus. This happens because the inner energy levels are filled with electrons, which shield the outer electrons from the pull of the nucleus and thus, the outer electrons remain far from the nucleus. Consequently, the size of the atom and hence the atomic radii increases.

Figure 3.4 (b) Variation of atomic radius with atomic number for alkali metals and halogens

• Down the group, the added electron enters the higher principal shell. As a result, the distance between the nucleus and outer electron increases. Hence, atomic radius increases down the group.

Vander Waal’s radius:

Definition box: Vander Waal’s radius: One half of the distance between the nuclei of two adjacent identical atoms belonging to two neighbouring molecules of an element in the solid state is known as Vander Waal’s radius.

• Vander Waal’s radius = XY/2
• The attractive force existing between molecules is Vander Waal’s force and hence, the name Vander Waal’s radius.
• The atomic radius that exists for noble gases is Vander Waal’s radius.

Metallic radius:

Definition box: Metallic radius is taken as half the intermolecular distance separating the metal cores in the metallic crystal.

Note box: Though an increase in nuclear charge increases the force of attraction on the electrons, the effect of adding a new shell is greater than the effect of inward pull because the electrons in the inner shells decrease the force of attraction of nucleus on the electrons in the outer shell. This is known as the screening effect. Hence, the atomic radius increases from top to bottom in a group. In transition metal series, Zr (Zirconium) and Hf (Hafnium) have comparable size due to lanthanide contraction.

(b) Ionic Radius

Definition box: Ionic radius is the distance from the centre of nucleus of an ion up to which it exerts its influence on its electron cloud. In other words, ionic radius is the distance between the nucleus and the outermost shell containing electron of an ion.

Radii of cations and anions:

• A cation (positive ion) is formed when one or more electrons are removed from the outer shell of a neutral atom

$$ \ce { Li -> Li^+ + e^- ; Mg -> Mg^{+2} + 2e^- ;Al -> Al^{+3} + 3e^- } $$

The removal of outer electron/s makes the number of positive charges in the nucleus greater than the number of electrons, i.e, the effective nuclear charge increases. Hence, the remaining electrons are strongly attracted by the nucleus and the electron cloud shrinks. This tends to decrease the radius of the cation. Therefore, a cation is smaller than the parent atom.

$$ \ce { Li -> Li^+ + e^- ; Na -> Na^+ + e^- ; Mg -> Mg^{+2} + 2e^- } $$
$$ 134 \space pm \hspace{10mm} 60 \space pm \hspace{10mm} 154 \space pm \hspace{10mm} 95 \space pm \hspace{10mm} 130 \space pm \hspace{10mm} 65 \space pm \hspace{10mm} $$

Higher the positive charge,smaller is the cation.Eg:- $Fe^{+3}$ is smaller than $Fe^{+2}$

• An anion (negative ion ) is formed when one or more electrons are added to the neutral atom

$$ \ce { F + e^- -> F^- ; Cl + e^- -> Cl^- } $$

The addition of electrons to the outer shell of an atom increases the net negative charge.The force of attraction on electrons decreases due to decrease in the effective nuclear charge .Hence, the electron cloud expands.Therefore , an anion is larger than the parent atom.

$$ \ce { F + e^- -> F^- ; Cl + e^- -> Cl^- } $$

$$ 72 \space pm \hspace{10mm} 136 \space pm \hspace{10mm} 99 \space pm \hspace{10mm} 181 \space pm $$

Higher the negative charge, larger is the anion.Eg: $O^{2-}$ is larger than $O^-$

• Sum of the radii of the two oppositely charged ions in an ionic solid is equal to interionic distance.

Interionic distance = $ r_{M^+} + r_{X^-} $

• Ions which have same number of electrons but different nuclear charges are known as isoelectronic  ions.

$Na^+,Mg^{2+},Al^{-3}$ are isoelectronic cations

$N^{3-},O^{2-},F^- $ are isoelectronic anions

• In case of isoelectronic ions, ionic radius decreases with increase in nuclear charge.
• $H^-$ and $I^-$ are the smallest and largest anions respectively.
• $H^-$ and $Cs^+$ are the smallest and largest cations respectively.

Example 3.5: Which of the following species will have the largest and the smallest size? Mg, $Mg^{2+}$, Al, $Al^{3+}$ .

Solution:

Atomic radii decrease across a period. Cations are smaller than their parent atoms. Among isoelectronic species, the one with the larger positive nuclear charge will have a smaller radius.Hence the largest species is Mg; the smallest one is Al$^{3+}$.

c) Ionization Enthalpy

Definition box: Ionization Enthalpy: Ionization Enthalpy of an element is defined as the minimum energy required to remove an electron from an isolated gaseous atom in its ground state. In other words, Ionization Enthalpy is the quantitative measure of the tendency of an element to lose electron.

• Ionization Enthalpy is depicted as $Δ_iH$.
• Since as $Δ_iH$.  is a measure of ease with which an electron can be removed, the smaller the value of as $Δ_iH$., easier is it to remove the electron from the atom.
• The ionization enthalpy is expressed in units of kJ $mol^{–1}$.
• We can define the first ionization enthalpy ($Δ_iH_1$) as the energy required to remove the first electron. The first ionization enthalpy for an element X is the enthalpy change ($Δ_iH_1$) for the reaction depicted in equation:           

$$ X(g) → X^+(g) + e^–  ….(3.1)$$

Figure 3.5 Variation of first ionization enthalpies ($Δ_iH$) with atomic number for elements with Z = 1 to 60

• We can define the second ionization enthalpy ($Δ_iH_2$) as the energy required to remove the second most loosely bound electron. It is the energy required to carry out the reaction shown in equation:

$$ X^+(g) → X^{2+}(g) + e^– …. (3.2)$$

• Energy is always required to remove electrons from an atom and hence ionization enthalpies are always positive. The second ionization enthalpy will be higher than the first ionization enthalpy because it is more difficult to remove an electron from a positively charged ion than from a neutral atom. In the same way the third ionization enthalpy will be higher than the second and so on as depicted:

           $$ Δ_iH_1 < Δ_iH_2 < Δ_iH_3  $$

• Lower $Δ_iH$ indicates greater reactivity of an element.

Figure 3.6(a) First ionization enthalpies ($Δ_iH$ )of elements of the second period as a function of atomic number (Z) and Fig. 3.6(b) $Δ_iH$  of alkali metals as a function of Z.

Factors influencing ionization enthalpy:

1) Atomic radius: Smaller the atomic radius, the higher is the ionization enthalpy. Thus, $Δ_iH$ increases across the period since the atomic size decreases and; decreases down the group since the atomic size increases.

2) Nuclear charge: the higher the positive charge of the nucleus, stronger is it’s attraction for the electrons, and hence the harder it is to remove the electrons. Thus, $Δ_iH$ increases across the period since the atomic size decreases and nuclear charge increases and; decreases down the group since the atomic size increases and nuclear force exerted on outermost electron decreases. Hence, the removal of an electron becomes easier.

3) Orbital penetration: It’s easier to remove electrons from p orbitals than from s orbitals, because the s orbitals penetrate towards the nucleus more closely than the p orbitals, thus making the electrons in the s orbital feel greater nuclear attraction. Thus, ionization enthalpy for an s-orbital is greater than that for a p-orbital.

4) Electron pairing: Within a subshell, paired electrons are easier to remove than unpaired ones. This is because repulsion between electrons in the same orbital is higher than repulsion between electrons in different orbitals. Thus, ionization enthalpy is greater for unpaired electrons than for paired.

5) Shielding or screening effect of the inner orbitals: Due to the electrons present in the inner orbitals, the electrons in the outermost orbitals feel lesser attraction for the nucleus than expected. In other words, the inner electron orbitals effectively screen or shield the outermost electrons from the nucleus, due to which the ionization enthalpy decreases.

6) Extra stability of completely filled and half-filled subshells:

• This explains the large value of $Δ_iH$  for beryllium (completely filled s-orbital) and nitrogen (with half-filled p-orbital). $Δ_iH$  of Be is greater than that of B and v if N is greater than that of oxygen.
• Noble gases have the highest value of $Δ_iH$.
• In general, atoms with outer electronic configurations: $ns^2$ or $ns^2np3^$ or $ns^2np^6$have very high value for ionization enthalpy as these configurations are of high stability.

Example 3.6: The first ionization enthalpy ($ \Delta_i$H ) values of the third period elements, Na, Mg and Si are respectively 496, 737 and 786 kJ mol$^{–1}4. Predict whether the first $ \Delta_i$ H value for Al will be more close to 575 or 760 kJ mol$^{–1}$ ? Justify your answer.

Solution: It will be more close to 575 kJ mol$^{–1}$. The value for Al should be lower than that of Mg because of effective shielding of 3p electrons from the nucleus by 3s-electrons.

Note box: Significance of Ionization enthalpy (IE): Lower IE represents greater metallic character, greater basic character and greater reducing power. Increase in IE is not significant in transition metal series because, the differentiating electron enters the inner orbital (d-orbital) with the outer configuration remaining the same.

(d) Electron Gain Enthalpy (or electron affinity):

Definition box: Electron Gain Enthalpy is defined as the energy change involved when an electron is added to an isolated gaseous atom.

$$ X(g) + e^– → X^–(g) …… (3.3) $$

• It indicates the tendency of an atom to accept an electron.
• Electron gain enthalpy can be positive or negative depending on the element.
• For many elements, energy is released when an electron is added to the atom and the electron gain enthalpy is negative. For example, group 17 elements (the halogens) have very high negative electron gain enthalpies because they can attain stable noble gas electronic configurations by picking up an electron.
• On the other hand, noble gases have large positive electron gain enthalpies because the electron has to enter the next higher principal quantum level leading to a very unstable electronic configuration.
• It may be noted that electron gain enthalpies have large negative values toward the upper right of the periodic table preceding the noble gases.
• It is denoted as $Δ_{eg}H$.
• It is expressed in kJ $mol^{–1}$.
• Greater the value of $Δ_{eg}H$, greater is the reactivity of the element.
• Electron gain enthalpy becomes more negative with increase in the atomic number across a period because, the effective nuclear charge increases from left to right across a period and consequently, it will be easier to add an electron to a smaller atom since the added electron on an average would be closer to the positively charged nucleus.
• Electron gain enthalpy becomes less negative as we go down a group because the size of the atom increases and the added electron would be farther from the nucleus. This is generally the case (Table 3.7).

• However, electron gain enthalpy of O or F is less negative than that of the succeeding element. This is because, when an electron is added to O or F, the added electron goes to the smaller n = 2 quantum level and suffers significant repulsion from the other electrons present in this level. For the n = 3 quantum level (S or Cl), the added electron occupies a larger region of space and the electron-electron repulsion is much less.

Example 3.7: Which of the following will have the most negative electron gain enthalpy and which the least negative?
P, S, Cl, F. Explain your answer.

Solution: Electron gain enthalpy generally becomes more negative across a period as we move from left to right. Within a group, electron gain enthalpy becomes less negative down a group. However, adding an electron to the 2p-orbital leads to greater repulsion than adding an electron to the larger 3p-orbital. Hence the element with most negative electron gain enthalpy is chlorine; the one with the least negative electron gain enthalpy is phosphorus.

Note box: Factors influencing electron affinity: 1. Nuclear Charge Greater the nuclear charge, greater will be the attraction for the incoming electron and as a result larger will be the value of electron affinity. 2. Atomic Size Larger the size of an atom, larger will be the distance between the nucleus and the incoming electron. Thus, smaller will be force of attraction felt by incoming electron and hence smaller will be the value of Electron affinity. 3. Electronic Configuration The more stable the configuration of an atom, lesser will be its tendency to accept an electron and hence lower will be its value of electron affinity.

(e) Electronegativity

Definition box: Electronegativity: The ability of an atom in a chemical compound to attract shared pair of electrons towards itself is called electronegativity.

• Electronegativity is not a measureable quantity. However, a number of numerical scales of electronegativity of elements like, Pauling scale, Mulliken-Jaffe scale, Allred-Rochow scale have been developed.
• Electronegativity increases across the period and decreases down the group. As the attraction between the outer (or valence) electrons and the nucleus increases as the atomic radius decreases in a period, the electronegativity also increases. For the same reason, the electronegativity values decrease with the increase in atomic radii down a group.

• Electronegativity is inversely related to the metallic properties of elements. Consequently, metals have low electronegativity and non-metals have high electronegativity. Thus, the increase in electronegativity across a period is accompanied by an increase in non-metallic properties (or decrease in metallic properties) of elements. Similarly, the decrease in electronegativity down a group is accompanied by a decrease in non-metallic properties (or increase in metallic properties) of elements.

Note box: Fluorine has the highest electronegativity (4.0) and caesium has the lowest (0.7). Electronegativity of hydrogen is 2.1.  This is the dividing line between metals and non-metals. Metals have EN < 2 and non-metals have EN > 2. Inert gases have zero electronegativity. Electronegativity values are useful for writing proper formulae of inorganic compounds. The less electronegative element is written first and the more electronegative is written latter. For example, a) $SO_2$ and not $O_2S$ b) $OF_2$ and not $F_2O$ c) $H_2O$ and not $OH_2$

Summary of trends in periodic properties of elements:

3.7.2 Periodic Trends in Chemical Properties

Valency (or) Oxidation state:

• Valency is the combining capacity of an element.
• Valency of an element depends on the number of electrons present in the outer shell (that is, number of valence electrons).
• The valency of representative elements (s and p block elements) is equal to the number of electrons in the outermost orbitals and is found as: “eight minus the number of outermost electrons as shown below”. Transition and inner transition (d and f block) elements exhibit variable valencies but usually, 2 and 3.
• The valency of elements across a period with respect to hydrogen and oxygen increases from 1 to 4 and then decreases to 0. This is illustrated in the table below:

• Valency of elements within a group exhibit same valency as the number of valence electrons is same.
• Elements of 18^th group have zero valency since they are chemically inert.

Example box:   In, $OF_2$ and $Na_2O$, the order of electronegativity of the three elements involved in these compounds is: F > O > Na.   1. Each of the atoms of fluorine, with outer electronic configuration 2$s^22p^5$, shares one electron with oxygen in the $OF_2$ molecule. Being highest electronegative element, fluorine is given oxidation state –1. Since there are two fluorine atoms in this molecule, oxygen with outer electronic configuration $2s^22p^4$ shares two electrons with fluorine atoms and thereby exhibits oxidation state: +2.   2. In $Na_2O$, oxygen being more electronegative accepts two electrons, one from each of the two sodium atoms and, thus, shows oxidation state –2. On the other hand sodium with electronic configuration $3s^1$ loses one electron to oxygen and is given oxidation state +1.

• Thus, the oxidation state or valency of an element in a particular compound can be defined as the charge acquired by its atom on the basis of electronegative consideration from other atoms in the molecule.

Example 3.8: Using the Periodic Table, predict the formulas of compounds which might be formed by the following pairs of elements; (a) silicon and bromine (b) aluminium and sulphur.

Solution:
(a) Silicon is group 14 element with a valence of 4; bromine belongs to the halogen family with a valence of 1. Hence the formula of the compound formed would be $SiBr_4$.

(b) Aluminium belongs to group 13 with a valence of 3; sulphur belongs to group 16 elements with a valence of 2. Hence, the formula of the compound formed would be $Al_2S_3$.

Some periodic trends observed in the valence of elements (hydrides and oxides) are shown in Table 3.9.

Anomalous Properties of Second Period Elements

• The first element of each of the groups 1 (lithium) and 2 (beryllium) and groups: 13-17 (boron to fluorine) differs in many respects from the other members of their respective group.
• For example, lithium, unlike other alkali metals, and beryllium unlike other alkaline earth metals, form compounds which are predominantly, covalent whereas, other elements of s-block form compounds that are predominantly, ionic in nature.
• In fact, the behaviour of lithium and beryllium is more similar with the second element of the following group i.e., magnesium and aluminium, respectively. This sort of similarity where the first element of each group resembles the second element of subsequent group is referred to as diagonal relationship.

• The reason for this anomalous behaviour is attributed to their small size, large charge, large radius ratio and high electronegativity of the elements. In addition, the first member of group has only four valence orbitals (2s and 2p) available for bonding, whereas the second member of the groups have nine valence orbitals (3s, 3p, 3d).
• Therefore, the maximum covanlency of the first member of each group is 4 (for example, boron can only form $[BF_4]^-)$ whereas, the other members of the groups can extend their valence shell to accommodate more than four pairs of valence electrons (for example, aluminium forms $[AlF_6]^{3-}$).

Example: Are the oxidation state and covalency of Al in $[AlCl(H_2O)_5]^{2+}$ same? 

Solution: No. The oxidation state of Al is +3 and the covalency is 6.

3.7.3 Periodic Trends and Chemical Reactivity

I. Reducing and oxiding nature:

According to electronic concept, oxidizing nature refers to the electron accepting tendency and reducing nature refers to electron releasing tendency of atoms. This implies that, metals are strong reducing agents and non-metals are strong oxidizing agents.

II. Nature of oxides:

• Metals form basic oxides while, non-metals form acidic or neutral oxides.
• Basic oxides dissolve in water to form hydroxides (bases) while, acidic oxides dissolve in water to form acids.

$$ \ce { Na_2O + H_2O -> 2NaOH } $$

$$ \ce { CO_2 + H_2O -> H_2CO_3}$$

Example: Show by a chemical reaction with water that $Na_2O$ is a basic oxide and $Cl_2O_7$ is an acidic oxide. 

Solution: 

$Na_2O$ with water forms a strong base whereas $Cl_2O_7$ forms strong acid. 

$$Na_2O + H_2O \rightarrow 2NaOH$$ 

$$Cl_2O_7 + H_2O \rightarrow 2HClO_4$$ 

Their basic or acidic nature can be qualitatively tested with litmus paper.

Questions from section 3.7: 

1. Describe the terms:

a) Electropositive

b) Electronegative

2. Explain the variation in the following trends, across the period and down the group:

a) Atomic radius

b) Ionization enthalpy

c) Electron gain enthalpy

d) Electronegativity

e) Valency

3. Define:

a) Vander Waal’s radius

b) Metallic radius

c) Ionic radius

d) Valency

4. Explain the factors influencing ionization enthalpy.

5. Explain the anomalous behaviour of Second Period Elements.

6. What do you mean by oxidizing nature of atoms?