Trends in Physical Properties

Trends in Physical Properties

Atomic Radius

Atomic radius is defined as the size of an atom, and it is determined by measuring the distance between the center of the nucleus and the outermost electron in an atom.

Factors Affecting Atomic Radii

1.     Number of Shells

2.     Effective Nuclear Charge (ZEff)

3.     Screening Effect

Number of Shells: As the number of shells increases, so do the atomic radii.

Example:

Effective Nuclear Charge: As the effective nuclear charge in an atom increases, the atomic radius decreases.

Example:

Screening Effect: More the screening effect of the electrons in the completely filled orbitals on the outer electrons, the more will be the atomic radius.

Example:

PERIODIC VARIATION OF ATOMIC RADII

1.                 On moving down the group the valence shells become far away from the nucleus and thus the atomic radius increases.

2.                 On moving along the period, the effective nuclear charge increases, and thus the electron cloud is attracted more strongly towards the nucleus resulting in the contraction of atomic radius.

 

Covalent radius

Covalent radius (for non-metals) and metallic radius (for metals) can be determined by the concept of atomic radius. It also depends on the number of protons and the attraction between the electrons and the protons. Depending on whether the atom is a metal or non-metal, there are three different types of atomic radii. They are

1.     Metallic Radius

2.     Covalent Radius

3.     Van der Waals’ Radius

Metallic radius

It is half the internuclear distance that separates the metal cores in a metallic crystal.

van der Waals Radius: It is one-half of the distance between the nuclei of two identical non-bonded isolated atoms.

Ionic Radius

Ions are formed either by gaining or losing electrons of an atom. Ionic radius is the distance between the centre of the nucleus and the outermost electron in an ion.

Ionic radius decreases as we move from left to right across a period in the periodic table and increases as we move from top to bottom down a group. Generally, anions have a higher ionic radius than cations.

For the same element

Isoelectronic species

These are ions of the different elements which have the same number of electrons but different magnitudes of the nuclear charge. The size of isoelectronic ions decreases with the increase in the nuclear charge.

Ionization Enthalpy

 

The amount of energy (work) required to remove an electron from the last orbit of an isolated (free) atom in gaseous state is known as ionisation potential or energy or better first ionisation potential of the element, i.e.,

     

1.     The amount of energies required to remove the subsequent electrons (2nd, 3rd, ...) from the monovalent gaseous cation of the element one after the other are collectively called successive ionisation energies. These are  designated as I.E1, I.E2, I.E3, I.E4 and so on. It may be noted that.  I.E4 > I.E3> I.E2> I.E1  (for a particular element)

IE is expressed in eV/atom or kcal mol–1or kJ mol–1

Note that eV atom–1 = 23.06 kcal mol–1 = 96.3 kJ mol–1

2.     In general, the first I.E. increases along the period from left to right. However there are some exceptions to the general trend –

I.E. decreases from elements of group 2 3.

I.E. decreases from elements of group 15 16.

 

3.     In a group of the periodic table, the ionisation energy decreases from top to bottom.

Factors Affecting Ionisation Energy

1.     Size of the Atom

2.     Effective Nuclear Charge

3.     Screening Effect

4.     Penetration Effect

5.     Electronic Configuration

Size of the Atom: As the size of the atom increases, ionisation energy decreases. This is because as size increases, the distance between the outermost electron and the nucleus increases. Therefore less energy is required to remove the electron from the atom.

Effective Nuclear Charge (ZEff): As the effective nuclear charge increases, the force of attraction between the nucleus and the outermost electron increases. Therefore, more energy is required to remove an electron from the atom.

Screening Effect: More the screening effect in an atom, the less the ionisation energy. This is because, as the inner electrons screen the outermost electrons from the nucleus, the force of attraction between the nucleus and the outermost electrons decreases. Therefore, less energy is required to remove an electron from the atom.

Penetration Effect: As the penetration of electrons in different orbitals increases, ionisation energy increases. The order of energy required to remove electrons from the orbitals in the same shell iss>p>d>f. s orbital is closer to the nucleus and is therefore more penetrated towards the nucleus. Therefore it is easier to remove electrons from p,d and f orbitals as compared to s orbital.

Electronic Configuration: According to Hund’s rule, the stability of half-filled and fully-filled degenerate orbitals is extremely high. Therefore, the removal of electrons from a half-filled or fully-filled orbital required more energy.

The order of IE for fully-filled, half-filled and partially filled orbitals is as follows.

First Ionisation Energy: It is the energy required to remove the outermost electron from a gaseous neutral atom in the ground state.

Example:

In the above example, the first ionisation energy of beryllium is greater than that of boron. This is because the outermost electron in the case of beryllium is present in the s-orbital which is more penetrated towards the nucleus. In the case of boron, the outermost electron is present in the p-orbital which is less penetrated than the s-orbital. Therefore, removing an electron from a s-orbital required more energy than removing an electron from a p-orbital.

Similarly, the first ionisation energy of nitrogen is greater than that of oxygen. This is because the 2p-orbital in nitrogen is half-filled, which gives it more stability than the partially filled 2p orbital in oxygen.

In the 13th group, the first ionisation energy of thallium is greater than that of gallium. This is due to the poor shielding of outer electrons by the inner d-electrons in thallium and the lanthanoid contraction in gallium.

Second Ionisation Energy: The energy required to remove an electron from a mono-positive isolated gaseous ion.

Third ionisation Energy: The energy required to remove the third most loosely bound electron.

For an isolated atom/ion, the removal of an electron from a cation is more difficult than the removal of an electron from a neutral atom.

Therefore,

Generally, the ionisation energy of non-metals is greater than the ionisation energy of metals. Also, noble gases have the highest ionisation energy value in a given period. Caesium has the lowest ionisation energy and is therefore used in photoelectric cells.

Electron shielding and effective nuclear charge

Electron shielding refers to the phenomenon in which electrons in an atom repel other electrons in the atom, thereby partially shielding them from the positive charge of the nucleus. This effect occurs because electrons have a negative charge, so they repel each other. As a result, an electron in an atom will be partially shielded by the other electrons that surround it, reducing its attraction to the nucleus.

 

Effective nuclear charge (Zeff) is the net positive charge experienced by an electron in an atom, taking into account the electron shielding effect. In other words, it is the actual nuclear charge that an electron in an atom feels after accounting for the shielding effect of other electrons. The effective nuclear charge is calculated by subtracting the number of inner-shell electrons from the total number of protons in the nucleus.

 

The shielding effect of inner-shell electrons on outer-shell electrons increases with increasing atomic number. This is because the number of inner-shell electrons also increases with increasing atomic number, resulting in greater electron shielding. The effective nuclear charge also increases with increasing atomic number, as the number of protons in the nucleus increases.

 

The concept of electron shielding and effective nuclear charge is important in understanding the trends in the periodic table, such as the size of atoms, ionization energy, and electron affinity. For example, as one moves from left to right across a period, the effective nuclear charge increases, resulting in a greater attraction between the nucleus and outer-shell electrons. This leads to a decrease in atomic size and an increase in ionization energy. Conversely, as one moves down a group, the effective nuclear charge remains relatively constant, while the number of inner-shell electrons increases, resulting in greater electron shielding. This leads to an increase in atomic size and a decrease in ionization energy.

 

Electron Gain Enthalpy

The change in enthalpy when an electron is added to a neutral gaseous atom to convert it into a gaseous anion is called electron gain enthalpy.

Electron gain enthalpy is positive when energy is absorbed, and negative when energy is released.

When the electron gain enthalpy is positive, energy is absorbed on the addition of an electron, i.e. ΔegH is positive. In this case, the addition of an electron makes the atom unstable.

When the electron gain enthalpy is negative, energy is released on the addition of an electron, i.e. ΔegH is negative. In this case, the addition of an electron makes the atom stable.

Electron gain enthalpy increases from left to right and decreases from top to bottom of the periodic table.

Factors Affecting Electron Gain Enthalpy:

1.     Atomic Size

2.     Effective Nuclear Charge (ZEff)

3.     Screening Effect

4.     Electronic Configuration

Atomic Size: As the size of the atom increases, the magnitude of electron gain enthalpy decreases. This is because, as the size of the atom increases, the force of attraction between the nucleus and the last shell which receives the incoming electron decreases. Therefore, electron gain enthalpy decreases.

Effective Nuclear Charge (ZEff): The greater the effective nuclear charge, the greater the tendency of the atom to attract the incoming electron towards itself. Therefore, electron gain enthalpy increases.

Screening Effect: More the screening effect of the electrons in the completely filled orbitals on the outer electrons, the more will be the electron gain enthalpy. This is because, as the screening effect increases, the force of attraction between the nucleus and the outermost electrons decreases. Therefore. Electron gain enthalpy also decreases.

Electronic Configuration: Elements with half-filled and fully-filled orbitals are more stable. Therefore, adding an electron to a stable electron configuration requires more energy. Hence, electron gain enthalpy has a high positive value.

Example:

In the example given above, neon and beryllium have a fully-filled outer orbital and are therefore more stable. Thus, the electron gain enthalpy in the case of neon and beryllium is the maximum. This is followed by nitrogen, which has a half-filled outer electron configuration.

The magnitude of electron gain enthalpies of group 3 elements is greater than the corresponding second period p-block elements.

The order of electron gain enthalpies of group 17 elements follows the order:

In the above order, chlorine has a higher electron gain affinity than fluorine owing to its larger size. Due to the smaller size of the fluorine atom, adding an additional electron will cause interelectronic repulsion and result in instability of the atom. The same is the case with oxygen in group 16.

Chlorine has the highest negative electron gain enthalpy in the periodic table. The electron gain enthalpies of noble gases are positive i.e. it results in an unstable electronic configuration.

Successive Electron Gain Enthalpy: ΔegH for the addition of a second electron to a neutral atom is positive. This is because interelectronic repulsion outweighs nuclear attraction.

Example:

Electronegativity

Electronegativity is the property of an atom in a molecule to attract the shared pair of electrons towards itself. The electronegativity of any atom is not constant, rather it is relative to the element to which it is bonded. Electronegativity helps in predicting the type of bond formed between two atoms. There are various scales of electronegativity. They are as follows:

The Pauling scale is the most commonly used scale of electronegativity.

1.     Mulliken scale: On the Mulliken scale, electronegativity X is taken as average of IE and EA, i.e.,

where IE and EA are expressed in electron volts

 

or where IE and EA are expressed in kJ/mol

or  where IE and EA are expressed in kcal/mol

 

b.                 Pauling scale: This is the most widely used scale and is based upon bond energy data. According to Pauling, the difference in electronegativity of two atoms A and B is given by the relationship as

where XA and XB are electronegativities of the atoms A and B respectively while.

where EA–B, EA–A and EB–B represent bond dissociation energies of the bonds A-B, A-A and B-B respectively. The Pauling and the Mulliken scales are related to each other by the relation,

Factors Affecting Electronegativity

1.     Atomic Size

2.     Effective Nuclear Charge (ZEff)

3.     Magnitude of Positive Charge on the Atom

Atomic Size: As the size of the atom increases, the force of attraction between the nucleus and the outermost electron decreases. Therefore, the electronegativity of the atom decreases.

Effective Nuclear Charge (ZEff): As the effective nuclear charge increases, the force of attraction between the nucleus and the outermost electron increases. Therefore, the electronegativity of the atom increases.

Magnitude of Positive Charge on the Atom: As the magnitude of the positive charge on the atom increases, the force of attraction between the nucleus and the outermost electron increases. Therefore, the electronegativity of the atom increases.

Electronegativity increases as we move from left to right across a period. This is because, on moving from left to right across a period, more electrons are added to the same shell, which increases the force of attraction between the nucleus and the outermost electrons.

Electronegativity decreases as we move down a group. This is because, as we move down a group, subsequent shells are added and therefore, the distance between the nucleus and the outermost electrons increases. Therefore, the force of attraction between the nucleus and the outermost electrons decreases and subsequently electronegativity decreases.

Fluorine has a higher electronegative whereas caesium has the least electronegativity value. Alkali metals have the lowest, and halogens have the highest electronegativity in their respective periods. It also varies among metals and non-metals. Non-metals are more electronegative than metals.

Periodic Trends in Chemical Properties

Periodicity of Valence or Oxidation states

The electrons present in the outermost shell of an atom are called valence electrons and the number of these electrons determine the valence or the valency of the atom.

The orbitals present in the valence shell are called valence orbitals.

The valence of an atom equal to either the number of valence electrons are equal to 8 minus the number of valence electron.

Transition and inner transition elements, exhibit variable valency due to the involvement of not only the valence electrons but the d or f electrons as well.

Variation along a period

As we move across a period from left to right, the number of valence electron increases from 1 to 8 .

The oxidation state of an element in a given compound may be defined as the charge acquired by its atom on the basis of electronegativity of the other atoms in the molecule.

Variations within a group

When we move down the group, the number of valence electrons remain the same, therefore all the elements in a group exhibit the same valence.

For Ex: All the elements of group 1 have valence one while all the elements of group 2 have valence of 2.

Noble gases present in group 18 are zerovalent i.e. their valance is zero since these elements are chemically inert.

Anomalous Properties of Second-Period elements

The first element of groups 1 and 2 and group 13 to 17 differ in many respect from the other members of their respective group.

Some elements on the 2nd period shows similarities with the elements of the 3rd period present diagonally to each other, though belonging to different group.

Similarities in properties of the elements placed diagonally to each other is called diagonal relationship.

The reason for the different chemical behaviour of the first member of a particular group of elements in the s and p block compared to the other members of the same group are

1) small size

2) large charge/ radius ratio

3) high electronegativity

4) absence of d orbital

5) the maximum covalency of first member of each group is 4 whereas other members of the groups can have a maximum covalency of 6.

6) because of the smaller size and high electronegativity, the first member of each group show a greater ability to form pπ-pπ multiple bonds either with itself or with other members of the second period.

Periodic Trends and Chemical Reactivity

The chemical and physical properties of the elements depend mainly upon their electronic configuration.

The atomic and ionic radii decrease as we move across the period from left to right. As a result ionization enthalpy generally increases and electron gain enthalpy becomes more negative across a period.

Ionization enthalpy of the extreme left element in a period is the lowest while electron gain enthalpy of the element at the extreme right is the most negative.

There is high chemical reactivity at the two extremes and the lowest in the centre.

The chemical reactivity of alkali metals on the extreme left is due to their ability to lose an electron to form corresponding cations.

Chemical reactivity of halogens on the extreme right is due to their ability to gain an extra electron leading to the formation of the anion.

Alkali metals are good reducing agents while halogens are good oxidizing agents.

Metallic character of an element which is highest at the extreme left decreases while the non-metallic character increases on moving across the period from left to right.

The elements at the extreme left of the periodic table readily combine with oxygen to form oxides which are most basic.

Elements on the extreme right from oxides which are most acidic.

Oxides of the elements in the centre are either amphoteric or neutral.

The amphoteric oxides show both acidic and basic properties.

Neutral oxides have neither acidic not basic properties.