Classification of Organic Halogen Compounds – Definition, Examples, Uses, & Facts

Classification of Organic Halogen Compounds – Definition, Examples, Uses, & Facts

A General Survey. Organic halogen compounds are derivatives of organic compounds in which one or more hydrogen atoms have been replaced by an equal number of halogen atoms (F, Cl, Br, or I). Almost any class of organic compounds (e.g., alcohols, ketones, carboxylic acids) can contain halogen atoms.

The haloalkanes, also known as alkyl halides, are a group of chemical compounds comprised of an alkane with one or more hydrogens replaced by a halogen atom (fluorine, chlorine, bromine, or iodine). The classification is determined by the number of carbons bonded to the carbon bearing the halide.

These are the compounds in which the halogen atom is bonded to an sp3-hybridised carbon atom next to carbon-carbon double bond (C=C) i.e. to an allylic carbon. These are the compounds in which the halogen atom is bonded to an sp3 hybridised carbon atom next to an aromatic ring.

Classification of Organic Halogen Compounds

Both chlorine and bromine are used as disinfectants for drinking water, swimming pools, fresh wounds, spas, dishes, and surfaces. They kill bacteria and other potentially harmful microorganisms through a process known as sterilization. Chlorine and bromine are also used in bleaching.

The halogens all form binary compounds with hydrogen, and these compounds are known as the hydrogen halides: hydrogen fluoride (HF), hydrogen chloride (HCl), hydrogen bromide (HBr), hydrogen iodide (HI), and hydrogen astatide (HAT). When in aqueous solution, the hydrogen halides are known as hydrohalic acids.

Classification of Organic Halogen Compounds img 1

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Environmental Pollution – Definition, History, Types, & Facts

Environmental Pollution – Definition, History, Types, & Facts

Any undesirable change in our environment that has harmful effects on plants, animals and human beings is called environmental pollution.

Environmental pollution is usually caused by the addition of waste products of human activity to the environment. The substances which cause pollution of environment are called pollutants. The pollutants may be solids, liquids or gaseous substances present in significant concentration in the environment.

Environmental Pollution

Our environment becomes polluted day by day, by the increased addition of industrial and domestic wastes to it. The air we breathe, the water we drink and the place where we live in, are highly contaminated.

The pollutants are classified as bio-degradable and non-biodegradable pollutants.

(i) Bio-Degradable Pollutants:

The pollutants which can be easily decomposed by the natural biological processes are called bio-degradable pollutants. Examples: plant wastes, animal wastes etc.

Environmental Pollution

(ii) Non Bio-Degradable Pollutants:

The pollutants which cannot be decomposed by the natural biological processes are called Non bio degradable pollutants. Examples: metal wastes (mainly Hg and Pb), D.D.T, plastics, nuclear wastes etc., These pollutants are harmful to living organisms even in low concentration. As they are not degraded naturally, it is difficult to eliminate them from our environment.

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Chemistry | Definition, Topics & History Concepts, Free Resources

Chemistry | Definition, Topics & History Concepts, Free Resources

Physical and Chemical Equilibrium 

Solutions

Chemical Bonding

Fundamentals of Organic Chemistry 

Basic Concepts of Organic Reactions  

Hydrocarbons

Haloalkanes and Haloarenes 

Environmental Chemistry 

Definition of Kirchhoff’s Rules and its Functions

Definition of Kirchhoff’s Rules and its Functions

Ohm’s law is useful only for simple circuits. For more complex circuits, Kirchhoff ’s rules can be used to find current and voltage. There are two generalized rules:

(i) Kirchhoff ’s current rule
(ii) Kirchhoff ’s voltage rule

Kirchhoff’s first rule (Current rule or Junction rule)

It states that the algebraic sum of the currents at any junction of a circuit is zero. It is a statement of law of conservation of electric charge. The charges that enter a given junction in a circuit must leave that junction since charge cannot build up or disappear at a junction. By convention, current entering the junction is taken as positive and current leaving the junction is taken as negative.

Definition of Kirchhoff's Rules and its Functions img 1

Applying this law to the junction A in Figure 2.23

I1 + I2 – I3 – I4 – I5 = 0
(or)
I1 + I2 = I3 + I4 + I5

Definition of Kirchhoff's Rules and its Functions

Example 2.20

For the given circuit find the value of I.

Definition of Kirchhoff's Rules and its Functions img 2

Solution

Applying Kirchhoff ’s rule to the point P in the circuit,
The arrows pointing towards P are positive and away from P are negative.
Therefore, 0.2A – 0.4A + 0.6A – 0.5A + 0.7A – I = 0
1.5A – 0.9A – I = 0
0.6A – I = 0
I = 0.6 A

Kirchhoff’s Second rule (Voltage rule or Loop rule)

It states that in a closed circuit the algebraic sum of the products of the current and resistance of each part of the circuit is equal to the total emf included in the circuit. This rule follows from the law of conservation of energy for an isolated system (The energy supplied by the emf sources is equal to the sum of the energy delivered to all resistors).

The product of current and resistance is taken as positive when the direction of the current is followed. Suppose if the direction of current is opposite to the direction of the loop, then product of current and voltage across the resistor is negative. It is shown in Figure 2.24 (a) and (b). The emf is considered positive when proceeding from the negative to the positive terminal of the cell. It is shown in Figure 2.24 (c) and (d)

Definition of Kirchhoff's Rules and its Functions img 3

Kirchhoff voltage rule has to be applied only when all currents in the circuit reach a steady state condition (the current in various branches are constant)

Definition of Kirchhoff's Rules and its Functions

Example 2.21

The following figure shows a complex network of conductors which can be divided into two closed loops like EACE and ABCA. Apply Kirchhoff’s voltage rule (KVR),

Definition of Kirchhoff's Rules and its Functions img 4

Solution

Thus applying Kirchhoff ’s second law to the closed loop EACE
I1R1 + I2R2 + I3R3 = ∈
and for the closed loop ABCA
I4R4 + I5R5-I2R2 = 0

Example 2.22

Calculate the current that flows in the 1 Ω resistor in the following circuit.

Definition of Kirchhoff's Rules and its Functions img 5

Solution

Definition of Kirchhoff's Rules and its Functions img 6

We can denote the current that flows from 9V battery as I1 and it splits up into I2 and (I1 – I2) at the junction E according Kirchhoff ’s current rule (KCR).

Now consider the loop EFCBE and apply KVR, we get

1I2 + 3I1 + 2I1 = 9
5I1 + I2 = 9 ………….. (1)

Applying KVR to the loop EADFE, we get
3(I1 – I2) – 1I2 = 6
3I1 – 4I2 = 6 ……….. (2)

Solving equation (1) and (2), we get
I1 = 1.83 A and I2 = – 0.13A

It implies that the current in the 1 ohm resistor flows from F to E.

Definition of Kirchhoff's Rules and its Functions

Wheatstone’s bridge

An important application of Kirchhoff ’s rules is the Wheatstone’s bridge. It is used to compare resistances and in determining the unknown resistance in electrical network. The bridge consists of four resistances P, Q, R and S connected as shown in Figure 2.25. A galvanometer G is connected between the points B and D. The battery is connected between the points A and C. The current through the galvanometer is IG and its resistance is G.

Definition of Kirchhoff's Rules and its Functions img 7

Applying Kirchhoff ’s current rule to junction B and D respectively

I1 – IG – I3 = 0 ……….. (2.45)
I2 + IG – I4 = 0 ……….. (2.46)

Applying Kirchhoff ’s voltage rule to loop ABDA,
I1P + IGG – I2R = 0 ………. (2.47)

Applying Kirchhoff’s voltage rule to loop ABCDA,
I1P + I3Q – I4S – I2R = 0 …………. (2.48)

When the points B and D are at the same potential, the bridge is said to be balanced. As there is no potential difference between B and D, no current flows through galvanometer (IG = 0). Substituting IG = 0 in equation (2.45), (2.46) and (2.47), we get

I1 = I3 …………. (2.49)
I2 = I4 …………… (2.50)
I1P = I2R …………. (2.51)

Using equation (2.51) in equation (2.48)
I3Q = I4S ………. (2.52)

Dividing equation (2.52) by equation (2.51), we get
\(\frac{P}{Q}\) = \(\frac{R}{S}\) …………. (2.53)

This is the condition for bridge balance. Only under this condition, galvanometer shows null deflection. Suppose we know the values of two adjacent resistances, the other two resistances can be compared. If three of the resistances are known, the value of unknown resistance (fourth one) can be determined.

Definition of Kirchhoff's Rules and its Functions

Example 2.23

In a Wheatstone’s bridge P = 100 Ω, Q = 1000 Ω and R = 40 Ω. If the galvanometer shows zero deflection, determine the value of S.

Solution

\(\frac{P}{Q}\) = \(\frac{R}{S}\)
S = \(\frac{Q}{P}\) × R
S = \(\frac{1000}{100}\) × 40
S = 400 Ω

Example 2.24

What is the value ofx when the Wheatstone’s network is balanced?

P = 500 Ω, Q = 800 Ω, R = x + 400, S = 1000 Ω

Definition of Kirchhoff's Rules and its Functions img 8

Solution

\(\frac{P}{Q}\) = \(\frac{R}{S}\), when the network is balanced
\(\frac{500}{800}\) = \(\frac{x+400}{1000}\)
x + 400 = \(\frac{5}{8}\) × 1000
x + 400 = 625
x = 625 – 400
x = 225 Ω

Definition of Kirchhoff's Rules and its Functions

Meter Bridge

The meter bridge is another form of Wheatstone’s bridge. It consists of a uniform wire of manganin AB of one meter length. This wire is stretched along a metre scale on a wooden board between two copper strips C and D. Between these two copper strips another copper strip E is mounted to enclose two gaps G1 and G2 as shown in Figure 2.26.

Definition of Kirchhoff's Rules and its Functions img 9

An unknown resistance P is connected in G1 and a standard resistance Q is connected in G2. A jockey (conducting wire-contact maker) is connected to the terminal E on the central copper strip through a galvanometer (G) and a high resistance (HR). The exact position of jockey on the wire can be read on the scale. A Lechlanche cell and a key (K) are connected between the ends of the bridge wire.

The position of the jockey on the wire is adjusted so that the galvanometer shows zero deflection. Let the position of jockey at the wire be at J. The resistances corresponding to AJ and JB of the bridge wire form the resistances R and S of the Wheatstone’s bridge. Then for the bridge balance

\(\frac{P}{Q}\) = \(\frac{R}{S}\) = \(\frac{r.AJ}{r.Jb}\) ……… (2.54)

where r is the resistance per unit length of wire.
\(\frac{P}{Q}\) = \(\frac{AJ}{JB}\) = \(\frac{l_{1}}{l_{2}}\) ……….. (2.55)
P = Q \(\frac{l_{1}}{l_{2}}\) ………… (2.56)

The bridge wire is soldered at the ends of the copper strips. Due to imperfect contact, some resistance might be introduced at the contact. These are called end resistances. This error can be eliminated, if another set of readings is taken with P and Q interchanged and the average value of P is found.

To find the specific resistance of the material of the wire in the coil P, the radius a and length l of the wire are measured. The specific resistance or resistivity ρ can be calculated using the relation.

Resistance = ρ \(\frac{l}{A}\)
By rearranging the above equation, we get
ρ = Resistance × \(\frac{A}{l}\) ………… (2.57)

If P is the unknown resistance equation (2.57) becomes,
ρ = P \(\frac{\pi a^{2}}{l}\)

Definition of Kirchhoff's Rules and its Functions

Example 2.25

In a meter bridge experiment with a standard resistance of 15 Ω in the right gap, the ratio of balancing length is 3:2. Find the value of the other resistance.

Solution

Q = 15 Ω, l1:l2 = 3:2
\(\frac{l_{1}}{l_{2}}\) = \(\frac{3}{2}\)
\(\frac{P}{Q}\) = \(\frac{l_{1}}{l_{2}}\)
P = Q \(\frac{l_{1}}{l_{2}}\)
P = 15 × \(\frac{3}{2}\) = 22.5 Ω

Example 2.26

In a meter bridge experiment, the value of resistance in the resistance box connected in the right gap is 10 Ω. The balancing length is l1 = 55 cm. Find the value of unknown resistance.

Solution

Q = 10 Ω

Definition of Kirchhoff's Rules and its Functions img 10

Potentiometer

Potentiometer is used for the accurate measurement of potential differences, current and resistances. It consists of ten meter long uniform wire of manganin or constantan stretched in parallel rows each of 1 meter length, on a wooden board. The two free ends A and B are brought to the same side and fixed to copper strips with binding screws. A meter scale is fixed parallel to the wire. A jockey is provided for making contact.

The principle of the potentiometer is illustrated in Figure 2.27. A steady current is maintained across the wire CD by a batteryBt.

Definition of Kirchhoff's Rules and its Functions img 11

The battery, key and the potentiometer wire connected in series form the primary circuit. The positive terminal of a primary cell of emf ε is connected to the point C and negative terminal is connected to the jockey through a galvanometer G and a high resistance HR. This forms the secondary circuit.

Let the contact be made at any point J on the wire by jockey. If the potential difference across CJ is equal to the emf of the cell ε, then no current will flow through the galvanometer and it will show zero deflection. CJ is the balancing length l. The potential difference across CJ is equal to Irl where I is the current flowing through the wire and r is the resistance per unit length of the wire.

Hence ε = Irl ………. (2.58)

Since I and r are constants, ε ∝ l. The emf of the cell is directly proportional to the balancing length.

Definition of Kirchhoff's Rules and its Functions

Comparison of emf of two cells with a potentiometer

To compare the emf of two cells, the circuit connections are made as shown in Figure 2.28. Potentiometer wire CD is connected to a battery Bt and a key K in series. This is the primary circuit.

The end C of the wire is connected to the terminal M of a DPDT (Double Pole Double Throw) switch and the other terminal N is connected to a jockey through a galvanometer G and a high resistance HR. The cells whose emf ε1 and ε2 to be compared are connected to the terminals M1, N1 and M2, N2 of the DPDT switch. The positive terminals of Bt, ε1 and ε2 should be connected to the same end C.

Definition of Kirchhoff's Rules and its Functions img 12

The DPDT switch is pressed towards M1, N1 so that cell ε1 is included in the secondary circuit and the balancing length l1 is found by adjusting the jockey for zero deflection. Then the second cell ε2 is included in the circuit and the balancing length l2 is determined. Let r be the resistance per unit length of the potentiometer wire and I be the current flowing through the wire.

We have ε1 = Irl1 ………. (2.59)
ε2 = Irl2 ………. (2.60)

By dividing equation (2.59) by (2.60)
\(\frac{\varepsilon_{1}}{\varepsilon_{2}}=\frac{l_{1}}{l_{2}}\) ……….. (2.61)

By including a rheostat (Rh) in the primary circuit, the experiment can be repeated several times by changing the current flowing through it.

Measurement of internal resistance of a cell by potentiometer

To measure the internal resistance of a cell, the circuit connections are made as shown in Figure 2.29. The end C of the potentiometer wire is connected to the positive terminal of the battery Bt and the negative terminal of the battery is connected to the end D through a key K1. This forms the primary circuit.

Definition of Kirchhoff's Rules and its Functions img 13

The positive terminal of the cell of emf ε whose internal resistance is to be determined is also connected to the end C of the wire. The negative terminal of the cell ε is connected to a jockey through a galvanometer and a high resistance. A resistance box R and key K2 are connected across the cell ε. With K2 open, the balancing point J is obtained and the balancing length CJ = l1 is measured. Since the cell is in open circuit, its emf is

ε ∝ l1 ………. (2.62)

A suitable resistance (say, 10 Ω) is included in the resistance box and key K2 is closed. Let r be the internal resistance of the cell. The current passing through the cell and the resistance R is given by

I = \(\frac{ε}{R+r}\)
The potential difference across R is
V = \(\frac{εR}{R+r}\)

Definition of Kirchhoff's Rules and its Functions

When this potential difference is balanced on the potentiometer wire, let l2 be the balancing length.

Then \(\frac{εR}{R+r}\) ∝ l2 ………… (2.63)
From equations (2.62) and (2.63)

Definition of Kirchhoff's Rules and its Functions img 14

Substituting the values of the R, l1 and l2, the internal resistance of the cell is determined. The experiment can be repeated for different values of R. It is found that the internal resistance of the cell is not constant but increases with increase of external resistance connected across its terminals.

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Electric Cells and Batteries

Electric Cells and Batteries

An electric cell converts chemical energy into electrical energy to produce electricity. It contains two electrodes (carbon and zinc) immersed in an electrolyte (sulphuric acid) as shown in Figure 2.17.

Electric Cells and Batteries img 1

Several electric cells connected together form a battery. When a cell or battery is connected to a circuit, electrons flow from the negative terminal to the positive terminal through the circuit.

By using chemical reactions, a battery produces potential difference across its terminals. This potential difference provides the energy to move the electrons through the circuit. Commercially available electric cells and batteries are shown in Figure 2.18

Electric Cells and Batteries img 2

Electric Cells and Batteries

Electromotive Force and Internal Resistance

A battery or cell is called a source of electromotive force (emf). The term ‘electromotive force’ is a misnomer since it does not really refer to a force but describes a potential difference in volts. The emf of a battery or cell is the voltage provided by the battery when no current flows in the external circuit. It is shown in Figure 2.19.

Electric Cells and Batteries img 3

Electromotive force determines the amount of work a battery or cell has to do move a certain amount of charge around the circuit. It is denoted by the symbol ε. An ideal battery has zero internal resistance and the potential difference (terminal voltage) across the battery equals to its emf.

In reality, the battery is made of electrodes and electrolyte, there is resistance to the flow of charges within the battery. This resistance is called internal resistance r. For a real battery, the terminal voltage is not equal to the emf of the battery. A freshly prepared cell has low internal resistance and it increases with ageing.

Determination of Internal Resistance

The circuit connections are made as shown in Figure 2.20.

Electric Cells and Batteries img 4

The emf of cell ε is measured by connecting a high resistance voltmeter across it without connecting the external resistance R as shown in Figure 2.20(a).

Electric Cells and Batteries img 5

Since the voltmeter draws very little current for deflection, the circuit may be considered as open. Hence the voltmeter reading gives the emf of the cell. Then, external resistance R is included in the circuit and current I is established in the circuit. The potential difference across R is equal to the potential difference across the cell (V) as shown in Figure 2.20(b).

Electric Cells and Batteries img 6

The potential drop across the resistor R is

V = IR ……….. (2.35)

Due to internal resistance r of the cell, the voltmeter reads a value V, which is less than the emf of cell ε. It is because, certain amount of voltage (Ir) has dropped across the internal resistance r.

Electric Cells and Batteries

Then V = ε – Ir
Ir = ε – V ………… (2.36)

Dividing equation (2.36) by equation (2.35), we get

\(\frac{Ir}{IR}\) = \(\frac{ε-V}{V}\)
r = [latex]\frac{ε-V}{V}[/latex]R ………… (2.37)

Since ε, V and R are known, internal resistance r can be determined. We can also find the total current that flows in the circuit.

Due to this internal resistance, the power delivered to the circuit is not equal to power rating mentioned in the battery. For a battery of emf ε, with an internal resistance r, the power delivered to the circuit of resistance R is given by

P = Iε = I (V + Ir) (from equation 2.36)

Here V is the voltage drop across the resistance R and it is equal to IR.
Therefore, P = I (IR + Ir)

P = I2R + I2r ……….. (2.38)

Here I2r is the power delivered to the internal resistance and I2R is the power delivered to the electrical device (here it is the resistance R). For a good battery, the internal resistance r is very small, then I2r << I2R and almost entire power is delivered to the external resistance.

Electric Cells and Batteries

Example 2.17

A battery has an emf of 12 V and connected to a resistor of 3 Ω. The current in the circuit is 3.93 A. Calculate (a) terminal voltage and the internal resistance of the battery (b) power delivered by the battery and power delivered to the resistor

Solution

The given values I = 3.93 A, ε = 12 V, R = 3 Ω

(a) The terminal voltage of the battery is equal to voltage drop across the resistor
V = IR = 3.93 × 3 = 11.79 V

The internal resistance of the battery
r = [latex]\frac{ε-V}{V}[/latex]R = [latex]\frac{12-11.79}{11.79}[/latex] × 3 = 0.05 Ω

(b) The power delivered by the battery P = Iε = 3.93 × 12 = 47.1 W
The power delivered to the resistor = I2R = 46.3 W

The remaining power P = (47.1 – 46.3) = 0.8 W is delivered to the internal resistance and cannot be used to do useful work. (It is equal to I2r).

Electric Cells and Batteries

Cells in Series

Several cells can be connected to form a battery. In series connection, the negative terminal of one cell is connected to the positive terminal of the second cell, the negative terminal of second cell is connected to the positive terminal of the third cell and so on. The free positive terminal of the first cell and the free negative terminal of the last cell become the terminals of the battery.

Suppose n cells, each of emf ε volts and internal resistance r ohms are connected in series with an external resistance R as shown in Figure 2.21

Electric Cells and Batteries img 7

The total emf of the battery = nε
The total resistance in the circuit = nr + R
By Ohm’s law, the current in the circuit is

Electric Cells and Batteries img 8

Case (a) If r<<R, then
I = \(\frac{nε}{R}\) ~ nI1 …….. (2.40)

where, I1 is the current due to a single cell

(I1 = \(\frac{ε}{R}\))

Thus, if r is negligible when compared to R the current supplied by the battery is n times that supplied by a single cell.

Case (b) If r>>R, I = \(\frac{nε}{nr}\) ~ \(\frac{ε}{r}\) …………. (2.41)

It is the current due to a single cell. That is, current due to the whole battery is the same as that due to a single cell and hence there is no advantage in connecting several cells.

Thus series connection of cells is advantageous only when the effective internal resistance of the cells is negligibly small compared with R.

Electric Cells and Batteries

Example 2.18

From the given circuit,

Electric Cells and Batteries img 9

Find

(i) Equivalent emf of the combination
(ii) Equivalent internal resistance
(iii) Total current
(iv) Potential difference across external resistance
(v) Potential difference across each cell

Solution

(i) Equivalent emf of the combination εeq = nε = 4 × 9 = 36 V

(ii) Equivalent internal resistance req = nr = 4 × 0.1 = 0.4 Ω

(iii) Total current I = \(\frac{nε}{R+nr}\)
= \(\frac{4×9}{10+(4×0.1)}\)
= \(\frac{4×9}{10+0.4}\) = \(\frac{36}{10.4}\)
I = 3.46 A

(iv) Potential difference across external resistance V = IR = 3.46 × 10 = 34.6 V. The remaining 1.4 V is dropped across the internal resistances of cells.

(v) Potential difference across each cell \(\frac{V}{n}\) = \(\frac{34.6}{4}\) = 8.65V

Electric Cells and Batteries

Cells in Parallel

In parallel connection all the positive terminals of the cells are connected to one point and all the negative terminals to a second point. These two points form the positive and negative terminals of the battery.

Let n cells be connected in parallel between the points A and B and a resistance R is connected between the points A and B as shown in Figure 2.22. Let ε be the emf and r the internal resistance of each cell.

Electric Cells and Batteries img 10

The equivalent internal resistance of the battery is \(\frac{1}{r_{c q}}\) = \(\frac{1}{r}\) + \(\frac{1}{r}\) + ………. \(\frac{1}{r}\)(n terms) = \(\frac{n}{r}\). So req = \(\frac{r}{n}\) and the total resistance in the circuit = R + \(\frac{r}{n}\). The total emf is the potential difference between the points A and B, which is equal to ε. The current in the circuit is given by

Electric Cells and Batteries img 11

where I1 is the current due to a single cell (\(\frac{ε}{R}\)) when R is negligible. Thus, the current through the external resistance due to the whole battery is n times the current due to a single cell.

Case (b)
If r<<R, I = \(\frac{ε}{R}\) ………… (2.44)

The above equation implies that current due to the whole battery is the same as that due to a single cell. Hence it is advantageous to connect cells in parallel when the external resistance is very small compared to the internal resistance of the cells.

Electric Cells and Batteries

Example 2.19

For the given circuit

Electric Cells and Batteries img 12

Find

(i) Equivalent emf
(ii) Equivalent internal resistance
(iii) Total current (I)
(iv) Potential difference across each cell
(v) Current from each cell

Solution

(i) Equivalent emf εeq = 5 V

(ii) Equivalent internal resistance,
Req = \(\frac{r}{n}\) = \(\frac{0.5}{4}\) = 0.125 Ω

(iii) total current, I = \(\frac{ε}{R+r/n}\)
I = \(\frac{5}{10+0.125}\) = \(\frac{5}{10.125}\)
I ~ 0.5 A

(iv) Potential difference across each cell
V = IR = 0.5 × 10 = 5V

(v) Current from each cell, I’ = \(\frac{I}{n}\)
I’ = \(\frac{0.5}{4}\) = 0.125 A

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Alkenes – Formula, Definition, Structure, Properties

Alkenes – Formula, Definition, Structure, Properties

Alkenes are unsaturated hydrocarbons that contain carbon-carbon double bond. They are represented by the general formulae CnH2n where ‘n’ stands for number of carbon atoms in the molecule. Alkenes are also known as olefins (in Latin oil maker) because the first member ethene combines with chlorine gas to form an oily liquid as a product.

(i) Nomenclature of Alkenes:

Let us write the IUPAC name for the below mentioned alkanes by applying the general rules of nomenclature that we already discussed in unit No.11

Alkenes:

IUPAC NAMES FOR ALKENES

Alkenes img 1

(ii) Isomerism:

Presence of double bond in Alkene provides the possibility of both structural and geometrical isomerism.

Structural Isomerism:

The first two member’s ethene C2H4 and propene C3H6 do not have isomers because the carbon atoms in the molecules can be arranged only one distinct way. However from the third member of alkene family butene C4H8, structural isomerism exists.

Alkenes img 2

structures (i) & (ii) are position isomers. structures (i) & (iii), (ii) & (iii) are chain isomers

Alkenes:

Geometrical Isomerism:

It is a type of stereoisomerism and it is also called cis-trans isomerism. Such type of isomerism results due to the restricted rotation of doubly bounded carbon atoms. If the similar groups lie on the same side, then the geometrical isomers are called Cis-isomers. When the similar groups lie on the opposite side, it is called a Trans isomer.

For Example:
The geometrical isomers of 2-Butene is expressed as follows

Alkenes img 3

General methods of preparation of alkenes:

(1) Preparation of alkene by dehydration of alcohol:

When an alcohol is heated at 430-440 K with excess of concentrated sulphuric acid, a molecule of water from alcohol is removed and an alkene is formed. This reaction is called elimination reaction.

Alkenes img 4

Ethene can also be prepared in laboratory by catalytic dehydration of alcohol.

Alkenes img 5

Alkenes:

(2) Preparation of alkenes from alkynes:

Alkynes can be reduced to cisalkenes using Lindlar’s catalyst (CaCO3 supported in palladuium partially deactivated with sulphur (or) gasoline). This reaction is stero specific giving only the cis-alkene.

Alkenes img 6

Alkynes can also be reduced to transalkenes using sodium in liquid ammonia. This reaction is stereospecific giving only the trans-alkene.

Alkenes img 7

(3) Preparation of alkenes by dehydrohalogenaton of halo alkanes.

Halo alkanes react with alcoholic KOH and eliminate hydrohalide resulting in the formation of alkene.

Alkenes img 8

(4) Preparation of alkenes from vicinal dihalogen derivative of alkanes or vicinal dihalides

The compound in which two halogen atoms are attached to adjacent carbon-atoms are called as vicinal dihalides. When vicinal dihalides are warmed with granulated zinc in methanol, they lose a molecule of ZnX2 to form an alkene.

Alkenes img 9

(5) Preparation of ethene by kolbe’s electrolytic method:

When an aqueous solution of potassium succinate is electrolyzed between two platinum electrodes, ethene is produced at the anode.

Alkenes img 10

At anode

Alkenes img 11

Alkenes:

Physical properities of alkenes:

The first three members (Ethene, Propene and Butene) are gases, next fourteen members are liquids and the higher alkenes are waxy solids. They are all colourless and odourless except ethene which has a sweet smell.

1. The melting and boiling point of alkenes increases along the homologous series. Like alkanes, straight chain alkenes have high boiling point compared to its isomeric branched alkenes.

2. Alkenes are slightly soluble in water but readily in organic solvents.

Chemical Properties of Alkenes:

Alkenes are more reactive than alkanes due to the presence of a double bond. The σ- bond is strong but the π- bond is weak. The typical reactions of alkenes involve addition of an electrophile across the double bonds proceeding through ionic mechanism. However addition reactions proceed through free-radical mechanism also. Ozonolysis and polymerization are some of the characteristic reactions of alkenes.

(i) Addition Reactions

(ii) Addition of hydrogen: (Hydrogenation of alkenes)

Hydrogen adds on to alkenes in the presence of a metal catalyst (Ni, Pd (or) Pt) to yield corresponding alkanes. This is known as catalytic hydrogenation. This process is of great importance in the manufacture of vanaspathi from vegetable oil. This helps to prevent rancidity of vegetable oils.

(ii) Addition of halogens: (Halogenation of alkenes)

When alkene is treated with halogens like chlorine or bromine, addition takes place rapidly and forms 1, 2 – dihalo alkane (or) vicinal dihalide

Alkenes img 12

Iodine reacts very slowly to form 1, 2 – diiodo alkane which are unstable and regenerate the original alkene by elimination of iodine.

Alkenes img 13

Alkenes:

TEST FOR ALKENE:

Alkenes img 14

Bromine in water is reddish brown colour. When small amount of bromine water is added to an alkene, the solution is decolourised as it forms dibromo compound. So, this is the characteristic test for unsaturated compounds.

Markovnikoff ’s Rule:

“When an unsymmetrical alkene reacts with hydrogen halide, the hydrogen adds to the carbon that has more number of hydrogen and halogen add to the carbon having fewer hydrogen”. This rule can also be stated as in the addition reaction of alkene / alkyne, the most electro negative part of the reagent adds on to the least hydrogen attached doubly bonded carbon.

(iii) Addition of water:- (Hydration of alkenes)

Normally, water does not react with alkenes. In the presence of concentrated sulphuric acid, alkenes react with water to form alcohols. This reaction follows carbocation mechanism and Markovnikof ’s rule.

Alkenes img 15

(iv) Addition of hydrohalides: (Hydrohalogenation of Alkenes)

Hydrogen halides (HCl, HBr and HI) add to alkene to yield alkyl halides. The order of reactivity of different hydrogen halides is HI>HBr>HCl. It is an example for electrophilic addition.

Alkenes:

(a) Addition of HBr to symmetrical alkene:

Addition of HBr to symmetrical alkene (similar groups are attached to double bond) yields alkyl halides (haloalkanes)

Alkenes img 16

(b) Addition HBr to unsymmetrical alkene:

In the addition of hydrogen halide to an unsymmetrical alkene, two products are obtained.

Alkenes img 17

Mechanism:

Consider addition of HBr to propene

Step: 1

Formation of electrophile:

In H-Br, Br is more electronegative than H. When bonded electron moves toward Br, polarity is developed and creates an electrophile H+ which attacks the double bond to form carbocation, as shown below.

Alkenes img 18

Step: 2

Secondary carbocation is more stable than primary carbocation and it predominates over a the primary carbocation.

Step: 3

The Brθ ion attacks the 2° carbocation to form 2-Bromo propane, as the major product.

Consider addition of HBr to 3-methyl-1-butene. Here the expected product according to markovnikoff ’s rule is 2-bromo-3-methyl butane but the actual major product is 2-Bromo-2-methyl butane. This is because, the secondary carbocation formed during the reaction rearranged to more stable tertiary carbocation. Attack of Br on this tertiary carbocation gives the major product 2-bromo-2-methyl butane.

Alkenes img 19

Alkenes:

Carbocation Rearrangement

Alkenes img 20

The addition of HBr to an alkene in the presence of organic peroxide, gives the anti Markovnikof ’s product. This effect is called peroxide effect.

Alkenes img 21

Mechanism:

The reaction proceeds via free radical mechanism.

Step: 1

The weak O-O single bond linkages of peroxides undergoes homolytic cleavage to generate free radical.

Alkenes img 22

Step: 2

The radicals abstracts a hydrogen from HBr thus generating bromine radical.

Alkenes img 23

Step: 3

The Bromine radical adds to the double bond in the way to form more stable alkyl free radical.

Alkenes img 24

Step: 4

Addition of HBr to secondary free radical

Alkenes img 25

Alkenes:

Addition of HBr to secondary free radical

The H-Cl bond is stronger (430.5 kJmol-1) than H-Br bond (363.7 kJmol-1), thus H-Cl is not cleaved by the free radical. Thus H-I bond is weaker (296.8 kJ mol-1), than H-Cl bond. Thus H-I bond breaks easily but iodine free radicals combine to form iodine molecules instead of adding to the double bond and hence peroxide effect is not observed in HCl & HI.

Kharasch Addition

Metal catalysed free radical addition of CXCl5 Compounds to alkene is called Kharash addition reaction.

(v) Addition of sulphuric acid to alkenes

Alkenes react with cold and concentrated sulphuric acid to form alkyl hydrogen sulphate accordance with Markownikof ‘s rule. Further hydrolysis yields alcohol.

Alkenes img 26

(2) Oxidation:

(i) With cold dilute alkaline KMnO4 solution (Baeyer’s Reagent)

Alkenes react with Baeyer’s reagent to form vicinal diols. The purple solution (Mn7+) becomes dark green (Mn6+), and then produces a dark brown precipitate (Mn4+).

Alkenes img 27

(ii) With acidified KMnO4 Solution:

Alkenes react with acidified KMnO4 solution and are oxidised to ketones or carboxylic acid depends on the substituent at the olefinic carbon atom. The purple solution becomes colourless. This is one of the test for unsaturation.

Alkenes img 28

Alkenes:

(iii) Ozonolysis:

Ozonolysis is a method of oxidative cleavage of alkenes or alkynes using ozone and forms two carbonyl compounds. Alkenes react with ozone to form Ozonide and it is cleaved by Zn/H2O to form smaller molecules. This reaction is often used to identify the structure of unknown alkene or alkyne by detecting the position of double or triple bond.

Alkenes img 29

(iv) Polymerisation:

A polymer is a large molecule formed by the combination of larger number of small molecules. The process in known as polymerisation. Alkenes undergo polymerisation at high temperature and pressure, in the presence of a catalyst.

For Example

Alkenes img 30

Recycling plastics

Extensive use of polymers clogs up landfills and polute the environment. Because of diversity of polymers in consumer products, recycling requires sorting the polymers into various sub-types, labels with codes and symbols, which are then recycled separately.

Table shows the codes and symbols used in recycling of ethene-based additionpolymers.

(Lower the number, greater the ease of recycling the material)

Alkenes img 31

Alkenes:

Uses of Alkenes

1. Alkenes find many diverse applications in industry. They are used as starting materials in the synthesis of alcohols, plastics, liquors, detergents and fuels.

2. Ethene is the most important organic feed stock in the polymer industry. E.g. PVC, Sarans and polyethylene. These polymer are used in the manufacture of floor tiles, shoe soles, synthetic fires, raincoats, pipes etc.,

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Haloarenes – Definition, Classification, Uses

Haloarenes – Definition, Classification, Uses

Haloarenes are the compounds in which the halogen is directly attached to the benzene ring.

Haloarenes img 1

Nomenclature of Haloarenes

In the IUPAC nomenclature, the halo arenes are named by adding prefix halo before the name of the aromatic hydrocarbon. For naming disubstituted arenes, the relative position of the substituent 1, 2; 1, 3 and 1, 4 are indicated by the prefixes ortho, meta and para, respectively.

For poly haloarenes the numbering should be done in such a way that the lowest possible number should be given to the substituents and the name of the halogens are arranged in alphabetic order. Nomenclature can be well understood from the following examples.

Haloarenes img 2

Haloarenes

Nature of C – X Bond in Haloarenes

In halo arenes the carbon atom is sp2 hybridised. The sp2 hybridised orbitals are shorter and holds the electron pair of bond more tightly.

Halogen atom contains P-orbital with lone pair of electrons which interacts with π-orbitals of benzene ring to form extended conjugated system of π-orbitals. The delocalisation of these electrons give double bond character to C – X bond. The resonance structure of halobenzene is given as

Haloarenes img 3

Due to this double bond character of C – X bond in haloarenes, the C – X bond is shorter in length and stronger than in halo alkanes.

Example

Haloarenes img 4

1. Direct Halogenation

Chlorobenzene is prepared by the direct chlorination of benzene in the presence of lewis acid catalyst like FeCl3

Haloarenes img 5

Haloarenes

2. From Benzene Diazonium Chloride

Chloro benzene is prepared by Sandmeyer reaction or Gattermann reaction using benzene diazonium chloride.

(i) Sandmeyer Reaction

When aqueous solution of benzene diazonium chloride is warmed with Cu2Cl2 in HCl gives chloro benzene

Haloarenes img 6

3. Preparation of Iodobenzene

Iodobenzene is prepared by warming benzene diazonium chloride with aqueous KI solution.

Haloarenes img 7

4. Preparation of Fluorobenzene

Fluoro benzene is prepared by treating benzenediazonium chloride with fluoro boric acid. This reaction produces diazonium fluoroborate which on heating produces flourobenzene. This reaction is called Balz – schiemann reaction.

Haloarenes img 8

5. Commercial Preparation of Chloro Benzene (Raschig Process)

Chloro benzene is commercially prepared by passing a mixture of benzene vapour, air and HCl over heated cupric chloride. This reaction is called Raschig process.

Haloarenes img 9

Haloarenes

Physical Properties

1. Melting and boiling points

The boiling points of monohalo benzene which are all liquids follow the order

Iodo > Bromo > Chloro

The boiling points of isomeric dihalobenzene are nearly the same

The melting point of para isomer is generally higher than the melting points of ortho and meta isomers. The higher melting point of p-isomer is due to its symmetry which leads to more close packing of its molecules in the crystal lattice and consequently strong intermolecular attractive force which requires more energy for melting.

p – Dihalo benzene > o – Dichloro benzene > m – Dichloro benzene

2. Solubility

Haloarenes are insoluble in water because they cannot form hydrogen bonds with water, but are soluble in organic solvents

3. Density

Halo arenes are all heavier than water and their densities follow the order.

Iodo benzene > Bromo benzene > Chloro benzene

Haloarenes

Chemical Properties

A. Reactions invoving halogen atom

1. Aromatic nucleophilic substitution reaction

Halo arenes do not undergo nucleophilic substitution reaction readily. This is due to C-X bond in aryl halide is short and strong and also the aromatic ring is a centre of high electron density.

The halogen of haloarenes can be substituted by OH, NH2, or CN with appropriate nucleophilic reagents at high temperature and pressure.

For Example

Haloarenes img 10

This reaction is known as Dow’s Process

Haloarenes img 11

2. Reaction with Metals

(a) Wurtz Fittig Reaction

Halo arenes reacts with halo alkanes when heated with sodium in ether solution to form alkyl benzene. This reaction is called Wurtz Fittig reaction

Haloarenes img 12

(b) Fittig reaction

Haloarenes react with sodium metal in dry ether, two aryl groups combine to give biaryl products. This reaction is called Fittig reaction

Haloarenes img 13

B. Reaction involving aromatic ring

3. Electrophilic substitution reaction

Haloarenes undergo aromatic electrophilic substitution reactions. The rate of eleclophilic substitution of halobenzene is lower than that of benzene halogen is deactivating due to – I effect of halogen. The lone pair of electrons on the chlorine involves in resonance with the ring. It increases the electron density at ortho and para position (refer figure no 14.1).

Haloarenes img 14

The halogen attached to the benzine ring with draw electron and thereby and hence the halogen which is attached to the benzene directs the incoming, electrophile either to ortho or to para position in electrophilie substitution reaction.

Toluene

Haloarenes

4. Reduction

Haloarenes on reduction with NiAl alloy in the presence of NaOH gives corresponding arenes.

Haloarenes img 15

5. Formation of Grignard Reagent

Haloarenes reacts with magnesium to form Grignard reagent in tetra hydrofuran (THF).

Haloarenes img 16

Uses of Chloro Benzene

  1. Chloro benzene is used in the manufacture of pesticides like DDT
  2. It is used as high boiling solvent in organic synthesis
  3. It is used as fire – swelling agent in textile processing

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OrganoMetallic Compounds – Definition, Details, Properties, and Applications

OrganoMetallic Compounds – Definition, Details, Properties, and Applications

Organo metallic compounds are organic compounds in which there is a direct carbon – metal bond. For Example CH3Mg I – Methyl magnesium iodide CH3CH2MgBr – Ethyl magnesium bromide

The Carbon – Magnesium bond in Grignard reagent is covalent but highly polar. The carbon atom is more electro negative than magnesium. Hence, the carbon atom has partial negative charge and the magnesium atom has partial positive charge

Organo Metallic Compounds img 1

Organo Metallic Compounds

Preparation

When a solution of alkyl halide in ether is allowed to stand over pieces of magnesium metal, the metal gradually dissolves and alkyl magnesium halide (Grignard reagent) is formed. All the reagents used should be pure and dry

Example

Organo Metallic Compounds img 2

Uses of Grignard Reagent

Grignard reagents are synthetically very useful compounds. These reagents are converted to various organic compounds like alcohols, carboxylic acids, aldehydes and ketones. The alkyl group being electron rich acts as a carbanion or a nucleophile. They would attack polarized molecules at a point of low electron density. The following reactions illustrate the synthetic uses of Grignard reagent.

1. Preparation of Primary Alcohol

Formaldehyde reacts with Grignard reagent to give addition products which on hydrolysis yields primary alcohol.

Organo Metallic Compounds img 3

2. Preparation of Secondary Alcohol

Aldehydes other than formaldehyde, react with Grignard reagent to give addition product which on hydrolysis yields secondary alcohol.

Organo Metallic Compounds img 4

Organo Metallic Compounds

3. Preparation of Tertiary Alcohol

Ketone reacts with Grignard reagent to give an addition product which on hydrolysis yields tertiary alcohols.

Example

Organo Metallic Compounds img 5

4. Preparation of Aldehyde

Ethyl formate reacts with Grignard reagent to form aldehyde. However, with excess of Grignard reagent it forms secondary alcohol

Example

Organo Metallic Compounds img 6

5. Preparation of Ketone

Acid chloride reacts with Grignard reagent to form ketones. However, with excess of Grignard reagent it forms tertiary alcohol.

Example

Organo Metallic Compounds img 7

6. Preparation of Carboxylic Acids

Solid carbon dioxide reacts with Grignard reagent to form addition product which on hydrolysis yields carboxylic acids.

For Example

Organo Metallic Compounds img 8

7. Preparation of Esters

Ethylchloroformate reacts with Grignard reagent to form esters.

Example

Organo Metallic Compounds img 9

Organo Metallic Compounds

8. Preparation of Higher Ethers

Lower halogenated ether reacts with Grignard reagent to form higher ethers.

Example

Organo Metallic Compounds img 10

9. Preparation of Alkyl Cyanide

Grignard reagent reacts with cyanogen chloride to from alkyl cyanide

Example

Organo Metallic Compounds img 11

Organo Metallic Compounds

10. Preparation of Alkanes

Compounds like water, alcohols and amines which contain active hydrogen atom react with Grignard reagents to form alkanes.

Example

Organo Metallic Compounds img 12

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VSEPR Theory – Postulates, Limitations, Predicting Shapes

VSEPR Theory – Postulates, Limitations, Predicting Shapes

Lewis concept of structure of molecules deals with the relative position of atoms in the molecules and sharing of electron pairs between them. However, we cannot predict the shape of the molecule using Lewis concept.

Lewis theory in combination with VSEPR theory will be useful in predicting the shape of molecules.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

Important Principles of VSEPR Theory are as follows:

1. The shape of the molecules depends on the number of valence shell electron pair around the central atom.

2. There are two types of electron pairs namely bond pairs and lone pairs. The bond pair of electrons are those shared between two atoms, while the lone pairs are the valence electron pairs that are not involved in bonding.

3. Each pair of valence electrons around the central atom repels each other and hence, they are located as far away as possible in three dimensional space to minimize the repulsion between them.

4. The repulsive interaction between the different types of electron pairs is in the following order.

lp – lp > lp – bp > bp – bp
lp – lone pair; bp – bond pair

The lone pair of electrons are localised only on the central atom and interacts with only one nucleus whereas the bond pairs are shared between two atoms and they interact with two nuclei. Because of this the lone pairs occupy more space and have greater repulsive power than the bond pairs in a molecule.

The following Table illustrates the shapes of molecules predicted by VSEPR theory. Consider a molecule ABx where A is the central atom and x represents the number of atoms of B covalently bonded to the central atom A. The lone pairs present in the atoms are denoted as L.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

Valence Shell Electron Pair Repulsion (VSEPR) Theory img 1

Valence Shell Electron Pair Repulsion (VSEPR) Theory img 2

Valence Shell Electron Pair Repulsion (VSEPR) Theory img 3

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Bond Parameters – Bond Order, Angle, Length, and Energy

Bond Parameters – Bond Order, Angle, Length, and Energy

A covalent bond is characterised by parameters such as bond length, bond angle, bond order etc. A brief description of some of the bond parameters is given below.

Bond Length

The distance between the nuclei of the two covalently bonded atoms is called bond length. Consider a covalent molecule A-B. The bond length is given by the sum of the radii of the bonded atoms (rA + rB). The length of a bond can be determined by spectroscopic, x-ray diffraction and electron-diffraction techniques. The bond length depends on the size of the atom and the number of bonds (multiplicity) between the combining atoms.

Bond Parameters img 1

Greater the size of the atom, greater will be the bond length. For example, carbon-carbon single bond length (1.54 Å) is longer than the carbon-nitrogen single bond length (1.43 Å). Increase in the number of bonds between the two atoms decreases the bond length. For example, the carbon-carbon single bond is longer than the carboncarbon double bond (1.33 Å) and the carbon-carbon triple bond (1.20 Å).

Bond Parameters

Bond Order

The number of bonds formed between the two bonded atoms in a molecule is called the bond order. In Lewis theory, the bond order is equal to the number of shared pair of electrons between the two bonded atoms. For example in hydrogen molecules, there is only one shared pair of electrons and hence, the bond order is one.  Similarly, in H2O, HCl, Methane, etc the central atom forms single bonds with bond order of one.

Bond Parameters img 2

Bond Angle

Covalent bonds are directional in nature and are oriented in specific directions in space. This directional nature creates a fixed angle between two covalent bonds in a molecule and this angle is termed as bond angle. It is usually expressed in degrees. The bond angle can be determined by spectroscopic methods and it can give some idea about the shape of the molecule.

Bond Parameters img 3

Bond Parameters

Bond Enthalpy

The bond enthalpy is defined as the minimum amount of energy required to break one mole of a particular bond in molecules in their gaseous state. The unit of bond enthalpy is kJ mol-1. Larger the bond enthalpy, stronger will be the bond.

The bond energy value depends on the size of the atoms and the number of bonds between the bonded atoms. Larger the size of the atom involved in the bond, lesser is the bond enthalpy.

In case of polyatomic molecules with, two or more same bond types, in the term average bond enthalpy is used. For such bonds, the arithmetic mean of the bond energy values of the same type of bonds is considered as average bond enthalpy. For example in water, there are two OH bonds present and the energy needed to break them are not same.

H2O(g) → H(g) + OH(g) ∆H1 = 502 kJ mol-1
OH(g) → H(g) + O(g) ∆H2 = 427 kJ mol-1

The average bond enthalpy of OH bond in water = \(\frac{502+427}{2}\) = 464.5 kJ mol-1

Bond Parameters img 4

Bond Parameters

Resonance

When we write Lewis structures for a molecule, more than one valid Lewis structures are possible in certain cases. For example let us consider the Lewis structure of carbonate ion [CO3]2-.

The skeletal structure of carbonate ion (The oxygen atoms are denoted as OA, OB & OC

Bond Parameters img 5

Total number of valence electrons = [1 × 4(carbon)] + [3 × 6 (oxygen)] + [2 (charge)]
= 24 electrons.
Distribution of these valence electrons gives us the following structure.

Bond Parameters img 6

Complete the octet for carbon by moving a lone pair from one of the oxygens (OA) and write the charge of the ion (2-) on the upper right side as shown in the figure.

Bond Parameters img 7

In this case, we can draw two additional Lewis structures by moving the lone pairs from the other two oxygens (OB and OC) thus creating three similar structures as shown below in which the relative position of the atoms are same.

They only differ in the position of bonding and lone pair of electrons. Such structures are called resonance structures (canonical structures) and this phenomenon is called resonance.

Bond Parameters img 8

It is evident from the experimental results that all carbon-oxygen bonds in carbonate ion are equivalent. The actual structure of the molecules is said to be the resonance hybrid, an average of these three resonance forms. It is important to note that carbonate ion does not change from one structure to another and vice versa.

It is not possible to picturise the resonance hybrid by drawing a single Lewis structure. However, the following structure gives a qualitative idea about the correct structure.

Bond Parameters img 9

It is found that the energy of the resonance hybrid (structure 4) is lower than that of all possible canonical structures (Structure 1, 2 & 3). The difference in energy between structure 1 or 2 or 3, (most stable canonical structure) and structure 4 (resonance hybrid) is called resonance energy.

Bond Parameters

Polarity of Bonds

Partial ionic character in covalent bond:

When a covalent bond is formed between two identical atoms (as in the case of H2, O2, Cl2 etc…) both atoms have equal tendency to attract the shared pair of electrons and hence the shared pair of electrons lies exactly in the middle of the nuclei of two atoms.

However, in the case of covalent bond formed between atoms having different electronegativities, the atom with higher electronegativity will have greater tendency to attract the shared pair of electrons more towards itself than the other atom. As a result the cloud of shared electron pair gets distorted.

Let us consider the covalent bond between hydrogen and fluorine in hydrogen fluoride. The electronegativities of hydrogen and fluorine on Pauling’s scale are 2.1 and 4 respective fluorine attracts the shared pair of electrons approximately twice as much as the hydrogen which leads to partial negative charge on fluorine and partial positive charge on hydrogen. Hence, the H-F bond is said to be polar covalent bond. Here, a very small, equal and opposite charges are separated by a small distance (91 pm) and is referred to as a dipole.

Dipole Moment:

The polarity of a covalent bond can be measured in terms of dipole moment which is defined as
μ = q × 2d

Where μ is the dipole moment, q is the charge and 2d is the distance between the two charges. The dipole moment is a vector and the direction of the dipole moment vector points from the negative charge to positive charge.

Bond Parameters img 10

The unit for dipole moment is columb meter (C m). It is usually expressed in Debye unit (D). The conversion factor is

1 Debye = 3.336 × 10-2

Diatomic molecules such as H2, O2 F2 etc have zero dipole moment and are called non polar molecules and molecules such as HF, HCl, CO, NO etc… have non zero dipole moments and are called polar molecules.

Molecules having polar bonds will not necessarily have a dipole moment. For example, the linear form of carbon dioxide has zero dipole moment, even though it has two polar bonds. In CO2, the dipole moments of two polar bonds (CO) are equal in magnitude but have opposite direction. Hence, the net dipole moment of the CO2 is, μ = μ1 + μ2 = μ1 + (-μ1) = 0.

Bond Parameters img 11

Incase of water net dipole moment is the vector sum of μ1 + μ2 as shown.

Bond Parameters img 12

Dipole moment in water is found to be 1.85D

Bond Parameters img 13

The extent of ionic character in a covalent bond can be related to the electro negativity difference to the bonded atoms. In a typical polar molecule, Aδ – Bδ+, the electronegativity difference (χA – xB) can be used to predict the percentage of ionic character as follows.

If the electronegativity difference (χA – χB), is equal to 1.7, then the bond A-B has 50% ionic character if it is greater than 1.7, then the bond A-B has more than 50% ionic character, and if it is lesser than 1.7, then the bond A-B has less than 50% ionic character.

Bond Parameters

Partial Covalent Character in Ionic Bonds:

Like the partial ionic character in covalent compounds, ionic compounds show partial covalent character. For example, the ionic compound, lithium chloride shows covalent character and is soluble in organic solvents such as ethanol.

The partial covalent character in ionic compounds can be explained on the basis of a phenomenon called polarisation. We know that in an ionic compound, there is an electrostatic attractive force between the cation and anion. The positively charged cation attracts the valence electrons of anion while repelling the nucleus.

This causes a distortion in the electron cloud of the anion and its electron density drifts towards the cation, which results in some sharing of the valence electrons between these ions. This, a partial covalent character is developed between them. This phenomenon is called polarisation.

The ability of a cation to polarise an anion is called its polarising ability and the tendency of an anion to get polarised is called its polarisability. The extent of polarisation in an ionic compound is given by the Fajans rules.

Fajans Rules

1. To show greater covalent character, both the cation and anion should have high charge on them. Higher the positive charge on the cation, greater will be the attraction on the electron cloud of the anion. Similarly higher the magnitude of negative charge on the anion, greater is its polarisability. Hence, the increase in charge on cation or in anion increases the covalent character.

Let us consider three ionic compounds aluminum chloride, magnesium chloride and sodium chloride. Since the charge of the cation increase in the order Na+ < Mg2+ < Al3+, the covalent character also follows the same order NaCl < MgCl2 < AlCl3.

2. The smaller cation and larger anion show greater covalent character due to the greater extent of polarisation. Lithium chloride is more covalent than sodium chloride. The size of Li+ is smaller than Na+ and hence the polarising power of Li+ is more. Lithium iodide is more covalent than lithium chloride as the size of I is larger than the Cl. Hence I will be more polarised than Cl by the cation, Li+.

3. Cations having ns2 np6 nd10 configuration show greater polarising power than the cations with ns2 np6 configuration. Hence, they show greater covalent character.

CuCl is more covalent than NaCl. Compared to Na+ (1.13 Å). Cu+ (0.6 Å) is small and have 3s2 3p6 3d10 configuration.

Electronic confiuration of Cu+
[Ar] 3d10
Electronic Confiuration of Na+
[He] 2s2, 2p6

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