Atomic Properties:

Properties such as atomic and ionic radii, ionization enthalpy, electron gain enthalpy, valency, screening effect, effective nuclear charge, and electronegativities are the properties of individual atoms. These properties are directly dependent on their electronic configuration.
Properties such as melting point, boiling point, the heat of fusion, heat of vaporization, density, atomic volume, etc., are collective properties of a group of atoms. These properties are indirectly dependent on their electronic configuration.
All such properties which are directly or indirectly dependent on the electronic configuration of the elements are called atomic properties.

Cause of Periodicity:

The physical and chemical properties of an element depend upon the distribution of electrons in the various shells of an atom. The distribution of electrons in the outermost shell is called the valence shell of an atom. This distribution of electrons in the outermost shell is important because it influences the physical and chemical properties of that element. In a particular group, all elements show a similar electronic distribution in the valence shell and hence they show similar physical and chemical properties. The following table gives the electronic distribution of the halogen family (Group – 17). We can observe the characteristic configuration of the ns2np5 valence shell. All these elements have similar properties with definite gradation (increase or decrease).

Thus the elements with similar configuration recur at regular intervals in the periodic table, the similar properties also recur in the periodic table. Thus the distribution of electrons in the outermost (valence) shell is the cause of periodicity.

In a period due to the gradual change in electronic configuration across the period from member to member, there is a gradual change in periodic property across the period. These properties include screening effect, atomic number, atomic radii, ionic radii, ionization enthalpy, electron gain enthalpy (electron affinity), electronegativity,  etc. They follow the general trend of periodicity.

In a group, valence shell configuration is the same, hence chemical properties in a group remain the same. There is a gradual change in physical properties due to an increase in atomic size which is due to the start of a new energy level.

Screening or Shielding Effect:

This effect is observed in an atom having more electrons and particularly more electron shells. The electrons in the valence shell are attracted by the positively charged nucleus. While there is repulsion between the valence electrons and the electrons present in the inner shells. Due to this, there is a decrease in the force of attraction between the electrons in the valence shell and the nucleus. This effect is known as the screening effect. The magnitude of the screening effect depends on the number of electrons in the inner shells.

The decrease in the force of attraction exerted by the nucleus on the valency electron due to the presence of electrons in the inner orbit is called screening effect or shielding effect.

Screening Effect Constant (Slater’s Rule):

Screening effect constant is denoted by letter σ. To find screening effect constant following steps should be followed.

Step 1: 

Write the electron configuration of the atom in the  form: (1s) (2s, 2p) (3s, 3p) (3d) (4s, 4p) (4d) (4f) (5s, 5p) . . .

Step 2: 

Identify the electron of interest (ns, np), and ignore all electrons in higher groups (to the right in the list from Step 1). These do not shield electrons in lower groups and hence do not contribute to screening effect constant.

Step 3: (For the shielding experienced by s- or p- electron):

All other electrons in the (ns, np) group contribute shielding to the extent of 0.35 each to the screening constant. note that for 1s this value is 0.30.

All the electrons in the (n – 1)th shell contribute 0.85 each to the screening effect constant.

All the electrons in the (n – 2)th shell contribute 1.0 each to the screening effect constant.

Step 3: (For the shielding experienced by d- or f- electron):

all other electrons in the (ns, np) group contribute shielding to the extent of 0.35 each to the screening constant. note that for 1s this value is 0.30.

All the electrons in a group lying left of (nd, nf) group contribute 1.0 each to the screening effect.

Examples of calculation of screening effect constant:

Screening Effect Down the Group (Alkali Metals):

We can see that as we move down the group screening effect increases.

Screening Effect Across the Period (Second Period):

We can see that as we move across the period from left to right there is increase in screening effect but there is no major changein the  screening effect as observed down the group in the periodic table.

Thus as the atomic number increases, the magnitude of the screening effect constant in case of s- and p- block elements increases in a period as well as in a group.

Calculation of screening effect constant for electron 3s orbital of bromine:

Atomic number of bromine is 35, its electronic configuration is 2, 8, 18, 7

The detailed configuration is  Br: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵

Br: (1s²)(2s²,2p⁶)(3s²,3p⁶)(3d¹⁰)(4s²,4p⁵)

Ignore the group to the right of the 3s electrons. These do not contribute to the shielding constant.

Screening effect constant = σ = 0.35 x 1 + 0.85 x 8 + 1 x 2 = 9.15

Calculation of screening effect constant for electron 3p orbital of bromine:

Ignore the group to the right of the 3p electrons. These do not contribute to the shielding constant.

Screening effect constant = σ = 0.35 x 5 + 0.35 x 2 + 0.85 x 8 + 1 x 2 = 11.25

Calculation of screening effect constant for electron 3d orbital of bromine: 

Ignore the group to the right of the 3d electrons. These do not contribute to the shielding constant.

Screening effect constant = σ = 0.35 x 9 + 1 x 8 + 1 x 8 + 1 x 2 = 21.15

Note: From above example we can see that for same main energy level screening effect constant for s- orbital is the least and for d- orbital it is the highest. So order for same main energy level is s < p < d< f.

Calculation of screening effect constant for electron 4s orbital of zinc: 

Atomic number of zinc is 30, its electronic configuration is 2, 8, 18, 7

The detailed configuration is  Zn 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰

Zn: (1s²)(2s²,2p⁶)(3s²,3p⁶)(3d¹⁰)(4s²)

Screening effect constant = σ = 0.35 x 1 + 0.85 x 18 + 1 x 8 + 1 x 2 = 25.65

Calculation of screening effect constant for electron 3d orbital of zinc: 

Br: (1s²)(2s²,2p⁶)(3s²,3p⁶)(3d¹⁰)(4s²)

Ignore the group to the right of the 3d electrons. These do not contribute to the shielding constant.

Screening effect constant = σ = 0.35 x 9 + 1 x 8 + 1 x 8 + 1 x 2 = 21.15

Note: The electrons in different orbitals are affected differently by the same nuclear charge dep[ending upon their proximity to the nucleus.

Effective Nuclear Charge:

Due to the screening effect, there is a decrease in the force of attraction on the electron in the valence shell towards the nucleus. Thus there is a decrease in the effect of nuclear charge. This reduced nuclear charge is called effective nuclear charge is denoted by  ‘Z_eff‘. The effective nuclear charge is the difference between the actual nuclear charge and the screening effect constant.charge.  Z_eff = Z – σ.

Let us consider the variation of effective nuclear charge across period:

Second Period Elements:

It is observed that the magnitude of effective nuclear charge increases in a period when we move from left to right.

Group – 1 (Alkali Metals):

Group – 2 (Alkaline Earth Metals):

It is observed that in a subgroup of normal elements the magnitude of effective nuclear charge remains almost the same.

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The word halogen is derived from the Greek words Halos (Sea Salt) and Genes (Producer). Since the binary compounds of halogen are the most abundant soluble salts found in the sea. Thus halogen means salt producers.

Electronic Configuration of Halogens:

Characteristic Electronic Configuration of Halogens:

• All halogens contain seven electrons in their outermost shell. All other shells are completely filled. They have characteristic outer orbit configuration of ns2 np5.
• The last electron during configuration occupies p orbital, hence these elements are p block elements.
• All orbits except the last orbit are completely filled. Hence they are normal elements.
• All halogens contain seven electrons in their outermost shell. Hence they are kept in group VII-A (17) of a periodic table, before inert gases.
• There are seven electrons in the outermost shell. So these elements require only one electron to complete the octet.

Properties of Halogens:

Monovalency of Halogens:

All halogen have shell electronic configuration is ns2 np5. They contain seven electrons in the valence shell. They need one electron to complete their octet. They attain the octet either by accepting an electron to form a univalent anion,  X-, (F-, Cl-, Br- and I-) by sharing the unpaired electron with the unpaired electron of another atom to form a covalent bond (as in Cl2, Br2, HCI, HF etc).  Therefore, the common valency of halogen family is 1. Hence they are monovalent.

Atomic Radius:

Atomic and ionic radius is the distance from the centre of the nucleus to the outermost shell of the atom or ion. From fluorine to iodine atomic radius increases because of following reasons.

Trend: As we go down in the group atomic radius increases from fluorine to iodine.

Explanation: As we go down in the group

• Atomic number and number of shell increases.
• Effective nuclear charge decreases.
• Screening effect increases.

It is to be noted that every ion is larger in size than the corresponding atom.

Oxidation State:

The characteristic electronic configuration of the halogens is ns2 np5. Hence, they tend to gain one electron to form the stable electronic configuration of the nearest noble gas atom and exhibit – 1 uniform oxidation state.

Chlorine, bromine and Iodine have empty n-‘d’ orbital. These elements when combining with the more electronegative element, their electrons of nth orbit get promoted to n-‘d’ orbital. Therefore, they can show positive oxidation states like +1, +3, +5 and +7.

Fluorine has only -1 oxidation state due to the absence of vacant n-’d’ orbitals.

Fluorine has Only -1 Oxidation State:

Fluorine is the most electronegative element having highest electronegativity 4. It does not have ‘d’ orbitals in Its valence shell. So it cannot expand the octet. Hence fluorine has only –1 oxidation state.

Ionization Potential (I.P.) or Ionization Enthalpy:

The energy required to remove outermost electron from the gaseous atom of an element, when it is in the ground state is called ionization potential or ionization enthalpy.

Since atomic radii of halogens are smallest in their respective period, their ionization potentials are very high. They have no tendency to lose the electron.

Among halogens, the I.P. (I.E.) value decreases with Increase in size of atom i.e. from fluorine to iodine. Therefore, the non-metallic properties decrease from fluorine to Iodine. Iodine is solid with metallic lustre.

Fluorine has Highest Ionization Enthalpy:

Ionisation potential is the minimum energy required to remove most loosely bound  electron from the outermost shell of an isolated gaseous atom. Fluorine has high. I.P. due to following reasons.

• It has the smallest atomic size
• The last electron present in 2p-orbital, which is nearer to the nucleus.

Therefore the last electron is held more tightly by the nucleus, due to greater nuclear charge. Thus more amount of energy is required to remove that electron and I.P. is more.

Electronegativity (EN):

The relative tendency of the bonded atom in a molecule to attract the shared electron pair towards itself is termed as its electronegativity.

Amongst the elements in the same period, halogens are most electronegative due to high nuclear charge and small atomic size. Their electronegativity and non-metallic character decrease gradually down the group with the increase in their atomic size. Electronegativity values are as follows F = 4,  Cl = 3, B r = 2. 8, 1 = 2. 5   (Note: Oxygen electronegativity is 3.51)

Fluorine Show High Electro-negativity:

Electro-negativity is the tendency of an atom to attract shared electrons towards itself in a molecule. Fluorine has high electro-negativity due to following reasons

• It has the smallest atomic size.
• The last orbit is second which is nearer to the nucleus and has a greater nuclear charge.

Therefore, the distance between the shared pair of electrons and nucleus of fluorine is small. Thus it has more ability to attract the shared pair of electrons towards itself and electro-negativity is more.

Electron Affinity (EA):

Electron affinity is the energy released when an electron is added to neutral gaseous atom forming a univalent negative ion. When halogens get electrons they give up energy.

Each halogen has maximum electron affinity in a period but in halogen family, it decreases from fluorine to iodine. The order of electron affinity is, Cl >  F> Br > I. It can be explained as follows

• Atomic size increases
• Effective nuclear charge decreases
• Ability to attract electron decreases
• Screening effect increases

Electron Affinity of Fluorine is less than that of Chlorine:

Electron affinity is the energy released when an electron is added to neutral gaseous atom forming a univalent negative ion. The electron affinity of fluorine is less than that of Chlorine. It can be explained as follows.

• Fluorine has an extremely small atomic size.
• 2p sub-shell is compact in fluorine.
• The maximum capacity of the outermost orbit in fluorine is eight electrons only.

Due to above reasons, the added electron comes too close to other valency electrons and this increases electron-electron repulsion. This results liberation of less energy when fluorine atom receives electron and forms F- Ions.

Bond Dissociation Energy (Bond Dissociation Enthalpy):

Bond dissociation energy is defined as the energy required to break a particular bond into atoms.

In general, halogens are diatomic molecules in which covalent bond is formed by overlapping of ‘p’ orbitals. Since atomic size increases from chlorine to Iodine, bond length increases from chlorine to iodine. As bond length increases from Cl to I. Bond dissociation energy decreases from chlorine to Iodine molecule.

Bond Dissociation Energy of Fluorine is Exceptionally Less:

Bond dissociation energy is the energy required to break the bond between the atoms in a gaseous molecule. Bond dissociation energy of F-F bond in F2 is less than Cl2, Br2 because of the following reasons

Fluorine has small atomic size and in fluorine, 2p subshell is compact and close to the nucleus.  Due to small atomic size number of electrons are held in a compact volume and there is strong repulsion amongst non bonded electrons. Hence bond becomes weak though bond is short.  Due to this reason bond dissociation energy of fluorine less than other halogens.

Multiple bonding takes place in Cl2, Br2, I2, due to the presence of d-orbitals, while such type of multiple bonding is absent in fluorine due to the absence of d-orbitals.  Hence bond dissociation energy is high for other halogens.  The order of bond dissociation energy is, Cl2 > Br2 > F2 > I2

Bond dissociation energy of fluorine molecule is much less than that of chlorine molecule.

This is because, in fluorine molecule, two fluorine atoms are quite close together (i.e. F-F bond length is very short) due to full size of fluorine atoms. This gives rise to very strong repulsion between the non-bonding electrons of the two atoms. The F-F bond, therefore, becomes weak.

In chlorine molecule, the repulsion is considerably minimised by longer Cl-Cl bond length due to the bigger size of the chlorine atoms.

General Properties of Halogens:

• State: Fluorine and Chlorine are gases, bromine is liquid and iodine is solid.
• Colour: All are coloured and the intensity of colour increases from fluorine to Iodine. Fluorine is light yellow/Pale yellow gas. Chlorine is greenish-yellow gas. Bromine is reddish-brown or Orange-red liquid and Iodine is violet or Shining black solid.
• Melting and Boiling Points: The elements have low boiling and melting points as their molecules are held together by Weak Vander Wall’s forces. These forces become stronger with an increase in atomic size. So boiling and melting points increase from fluorine to Iodine. Thus, fluorine and chlorine are gases, bromine is volatile liquid and Iodine is solid, which easily sublimes.
• Oxidizing Property: Halogens are strong oxidizing agents. Fluorine is the most powerful oxidizing agent. Oxidizing power decreases from fluorine to iodine.

Fluorine is the most powerful oxidizing agent:

• Fluorine has smallest atomic size.
• Fluorine has the highest electronegativity (= 4)
• F-F bond dissociation energy is very low.
• Fluorine atom accepts electron readily to form F- ion. Hence it has high reduction potential Eo Red = +2.87 Volt.

Reasons for Anomalous Behaviour of Fluorine:

Fluorine differs from other halogens due to following reasons

• Smallest atomic size.
• Highest electronegativity.
• Weak F-F bond i.e. F—F bond dissociation energy is low.
• Non-availability of Id’ orbitals in its valence shell.
• Strongest oxidizing agent (stronger than oxygen)

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