Oxidation Number OR Oxidation State:

The donation of electrons is called the oxidation and the gain of electrons is called the reduction. Oxidation and reduction can further be explained by a knowledge of “Oxidation number”.

The oxidation state of an atom in its free or ground state is taken as zero. When the atom loses electrons its oxidation state increases and when the atom gains electrons its oxidation state decreases. The term oxidation-number represents the positive or negative character of the atom in a compound.

Oxidation number is defined as the charge an atom appears to have when electrons are assigned in accordance with the following arbitrary rules.

Electrons shared by two like atoms are divided equally between the two atoms. Electrons shared between two unlike atoms are assigned to the more electronegative atom of them.

Conventions Used in Assigning Oxidation Number or Oxidation State:

• The oxidation number of an element in a free atomic state (Na, H, Cl, O, P etc) or in its poly-atomic state (graphite, H₂, Cl₂, O₂ etc) is always zero.
• The oxidation number of hydrogen is always +1 in its compounds.  However, in metal hydrides like NaH, MgH₂  etc. the oxidation number of hydrogen is -1 because metals are more electropositive than hydrogen.
• O.N. of oxygen is always -2 in its compounds.  However, in peroxides like H₂O₂, Na₂O₂, BaO₂ etc. the oxidation number of oxygen is -1. In OF the oxidation number of oxygen is +2 because F is more electronegative than O.
• O.N. of group IA element i.e. Li, Na, K etc is always +1 in their compounds.
• O. N. of group IIA elements i.e. Be, Mg, Ca, Sr and Ba are always +2 in their compounds.
• O. N. of F is always -1 in its compounds because it is most highly electronegative.  Oxidation O. N. of other elements of group VIIA. (17) i.e. Cl, Br and I are also generally –1.
• In an ion, the sum of the oxidation numbers of different atoms is equal to charge over the ion.
• In a complex compound (involving co-ordination by ligands) it is more convenient to use oxidation number of group (ligand) as a whole instead of the oxidation number of individual atoms. For example, in HCN the oxidation number of CN- ion is –1. Here CN-  as a whole is considered and not of individual C or N.
• on the basis of the above standard oxidation numbers, which may be taken as rules, the oxidation, a number of a particular given atom in a compound can be determined.

Valency and Oxidation State:

Valency is a different term than oxidation number though sometimes the valency and the oxidation number of an element are same in a compound. Valency of an element is given by the number of electrons it actually loses or gains or shares during the formation of a compound, Whereas oxidation number is just the apparent charge (not necessarily actual) over the atom when the electrons are counted according to the arbitrary rules given earlier.

In most of the cases, the valency of an element is constant whereas the oxidation state of an element may vary in its different compounds. Valency and oxidation states of carbon in its different compounds give a good example of this. In CH₄, CH₃Cl, CH₂Cl₂, CHCl₃ and CCl₄ the valency of carbon is always four (due to sharing of four electrons) but its oxidation number is – 4, -2, 0, +2 and +4 respectively.

Oxidation-Reduction in Terms of Oxidation Number:

On the basis of oxidation number a reaction involving the increase in oxidation number is called as oxidation while a reaction involving the decrease in oxidation number is called as reduction (Remember increase in O.N. means increase in positive O.N. or decrease in negative O.N., while decrease in O.N. means decrease in positive O.N. or increase in negative O.N.).

For example, in the reaction, 2Mg + O₂ →  2MgO, The O.N. of Mg increase from 0 to +2.while the O.N. of O decreases from 0 to -2.  Thus, magnesium is oxidised while oxygen is reduced.

Science > Chemistry > Redox Reactions > Oxidation Number or Oxidation State

The post Oxidation Number or Oxidation State appeared first on The Fact Factor.

In this article we shall study about redox reactions, in which both the oxidation and reduction reactions take place simultaneously.

Oxidation:

Old Concept:

• It is a process in which addition of oxygen takes place.

2Mg + O₂  →      2MgO

• It is a process in which addition of electronegative radical takes place.

2FeCl₂ +  Cl₂ →  2FeCl₃

• It is a process in which removal of hydrogen takes place.

H₂S + 2 [Cl] →  S + 2HCl

• It is a process in which removal of electropositive radical takes place.

2KI + H₂O₂ → I₂ + 2KOH

Moden Concept:

According to the electronic concept, a reaction in which loss of electrons from an atom or an ion takes place is called oxidation. Consider reaction, 

2Mg + O₂ →  2MgO

In this reaction, the valency of magnesium is increased’ from zero (in the atomic state) to + 2 (in MgO).

i.e.  Mg⁰ →  Mg²⁺ + 2 e^–

In this reaction, magnesium is losing electrons. And hence oxidation of magnesium takes place.

Reduction:

Old Concept:

• It is a process in which addition of hydrogen takes place.

Cl₂ + H₂  →  2HCl

• It is a process in which addition of electropositive radical takes place.

2HgCl₂ +  SnCl₂ → Hg₂Cl₂ + SnCl₄ .

• It is a process in which removal of oxygen takes place.

CuO + 2 [H] → Cu + 2H₂O

• It is a process in which removal of electronegative radical takes place.

FeCl₃ + H → FeCl₂ + HCl

Modern Concept:

According to the electronic concept, a reaction in which the gain of electrons by an atom or an ion takes place is called reduction. Consider reaction, 

2HgCl₂ +  SnCl₂ → Hg₂Cl₂ + SnCl₄

In this reaction, the valency of mercury is decreased’ from +2  (in HgCl₂) to +1 (in Hg₂Cl₂).

i.e.   Hg²⁺ +   e^–  →   Hg ⁺

In this reaction mercury is gaining electron. And hence reduction of mercury takes place.

Redox Reaction:

In any of a chemical reaction if one of the reactants is oxidized, other is surely reduced, i.e. oxidation and reduction always take place simultaneously.

Example – 1: 

2Mg + O₂ →  2MgO  

Mg is oxidized to MgO (addition of oxygen, i.e. increase in positive valency of Mg i.e. loss of electrons by Mg), whereas oxygen is reduced to MgO (addition of positive radical, i.e. increase in negative valency of oxygen, i.e. gain of electrons by oxygen)

Example – 2: 

2HgCl₂ +  SnCl₂ → Hg₂Cl₂ + SnCl₄

HgCl₂ is reduced to Hg₂Cl₂ whereas SnCl₂ is oxidised to SnCl₄.

Thus oxidation and reduction take place simultaneously.  Therefore, all such reactions are called as reduction-oxidation reactions or redox reactions.  In all such reactions, one of the reactants loses the electrons (oxidized) while other gains those electrons (reduced)

However, it should be remembered that all the chemical reactions are not redox reactions.  There are several other types of reactions also.

NaCl + AgNO₃ → AgCl + NaNO₃

In such reactions none of’ the reactants is oxidized or reduced; simply the exchange of cation or anion takes place.

Oxidizing Agent (Oxidant):

The substance which excepts electrons and makes the other substance to lose electrons is called oxidizing agent or oxidant.

Example – 1:

Consider reaction, 

2Mg + O₂ →  2MgO

In this reaction, oxygen is making magnesium to lose electrons and hence in this reaction oxygen is the oxidizing agent.

Example – 2:

Consider reaction,

2K + Cl₂  → 2KCl

In this reaction, chlorine is making potassium to lose an electron and hence in this reaction chlorine is the oxidizing agent.

Characteristics of Oxidizing Agent:

• The substance which excepts electrons and makes the other substance to lose electrons is called oxidizing agent or oxidant.
• In a reaction, the oxidizing agent oxidizes the other substance but is itself reduced.
• Oxygen, or a substance capable of giving oxygen, is always a good oxidizing agent.
• According to electron concept, an oxidizing agent is that which is capable of de-electronating the other substance.
• An oxidizing agent is an electron acceptor and during the redox reaction, it is electronated.
• Fluorine (F) has a maximum tendency to accept electrons hence it is the strongest oxidizing agent.

Examples of Common oxidizing Agents:

Oxygen (O or O₂), Ozone (O₃), Hydrogen peroxide (H₂O₂), Sulphuric acid (H₂SO₄), Nitric acid (HNO₃), Perchloric acid (HClO₄), Potassium chlorate (KClO₃), Acidified potassium dichromate (K₂ Cr₂ O₇ + H₂SO₄), Acidified potassium permanganate (KMnO₄ + H₂SO₄), Alkaline potassium permanganate (KMnO₄ + KOH)

Reducing Agent (Reductant):

The substance which loses electrons and makes the other substance to gain electrons is called reducing agent or reductant.

Example – 1:

Consider reaction, 

2Mg + O₂ →  2MgO

In this reaction, magnesium is making oxygen to gain electrons and hence in this reaction magnesium is reducing agent.

Example – 02:

Consider reaction, 2K + Cl₂  → 2KCl, In this reaction potassium, is making chlorine to gain an electron and hence in this reaction potassium is reducing agent.

Characteristics of Reducing Agent:

• The substance which loses electrons and makes the other substance to gain electrons is called reducing agent or reductant.
• In a reaction, the reducing agent reduces the other substance but is itself oxidised.
• Hydrogen, or a substance capable of giving hydrogen, is always a good reducing agent.
• According to electron concept, a reducing agent is that which is capable of electronating the other substance.
• A reducing agent is an electron donor and during the redox reaction, it is de-electronated.
• Sodium (Na) has a maximum tendency to donate electron hence it is the strongest reducing agent.

Examples of Common Reducing Agents:

Hydrogen (H or H₂), Hydrogen iodide (HI), Hydrogen sulphide (H₂S), Lithium aluminium hydride (LiAI H₄), Sodium borohydride (NaB H₄), Sulphur dioxide (SO₂), Carbon (C), Ozone (O₃), Hydrogen peroxide (H₂O₂), Tin & hydrochloric acid (Sn  + HCl), Sodium & alcohol (Na + C₂ H₅OH), Metallic salts (ous) like SnCl₂ , FeSO₄ etc.

Science > Chemistry > Redox Reactions > Introduction to Redox Reactions

The post Introduction to Redox Reactions appeared first on The Fact Factor.

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)

The post Halogens appeared first on The Fact Factor.