Melting Point, Boiling point, Solubility

Boiling point

            The boiling point of an element or a substance is the temperature at which the vapor pressure of the liquid equals the environmental pressure surrounding the liquid.

             The boiling point of a liquid varies dependent upon the surrounding environmental pressure. Different liquids (at a given pressure) boil at different temperatures. A liquid in a vacuum environment has a lower boiling point than when the liquid is at atmospheric pressure. A liquid in a high pressure environment has a higher boiling point than when the liquid is at atmospheric pressure.

            The normal boiling point of a liquid is the special case in which the vapor pressure of the liquid equals the defined atmospheric pressure at sea level, 1 atmosphere. At that temperature, the vapor pressure of the liquid becomes sufficient to overcome atmospheric pressure and lift the liquid to form bubbles inside the bulk of the liquid. The standard boiling point is now defined by IUPAC as the temperature at which boiling occurs under a pressure of 1 bar.

            The heat of vaporization is the amount of energy required to convert or vaporize a saturated liquid (i.e., a liquid at its boiling point) into a vapor.

            Liquids may change to a vapor at temperatures below their boiling points through the process of evaporation. Evaporation is a surface phenomenon in which molecules located near the liquid's edge, not contained by enough liquid pressure on that side, escape into the surroundings as vapor. On the other hand, boiling is a process in which molecules anywhere in the liquid escape, resulting in the formation of vapor bubbles within the liquid.

Determination of Boiling Points

·  Boiling point is determined by Capillary tube method.

·  In this method, a few drops of  liquid are placed in a thin walled small test tube.

·  A capillary tube sealed at about 1 cm from one end, is dropped in to it.

·  A glass tube containing the liquid and capillary, is then tied along a side of thermometer so that the liquid stands just near the bulb.

·  The thermometer is then lowered in a beaker containing paraffin oil.

·  The beaker is heated and the bath liquid stirred continuously using with stirrer.

·  When the boiling point reached, bubbles come from lower end of capillary.

·  The read the temperature from thermometer when the evaluation of bubbles just stop.

Melting Point

            The melting point of a substance is the temperature at which the solid phase converts to the liquid phase under 1 atmosphere of pressure.

            The melting point is one of a number of physical properties of a substance that is useful for characterizing and identifying the substance.

            To measure the melting point of a substance, it is necessary somehow to gradually heat a small sample of the substance while monitoring its temperature with a thermometer. The temperature at which liquid is first seen is the lower end of the melting point range. The temperature at which the last solid disappears is the upper end of the melting point range. A pure substance normally has a melting point range no larger than 1-1.5 ^oC.

            Although many substances melt cleanly and can be melted, crystallized, and remelted repeatedly without chemical decomposition, others chemically decompose before they melt, forming substances of lower molecular weight.

            The temperature at which the color change is first observed signals that the substance is approaching the decomposition temperature.

Determination of Melting Points

·  Melting point is also determined by Capillary tube method.

· A glass capillary tube which is 5-6 cm long and 1mm diameter, normally used to contain the sample for a melting point determination.

·  The tube must have one open end into which the sample can be loaded, and one sealed end so that the capillary will retain the solid sample.

·  The substance should stand in the capillary 3-4 mm from the bottom when thoroughly packed.

· The capillary is wetted with liquid in the bath and then tied along a side of thermometer fixed in an iron stand.

· The thermometer is then lowered in a beaker containing paraffin oil.

· The beaker is heated and the bath liquid stirred continuously using with stirrer.

·  When the substance in the capillary just shows sign of melting , the burner is removed and stirring continued.

· The read the temperature from thermometer when the substance melts and become transparent. This is the melting point range of that substance.

Solubility

Ø  A solution is a homogeneous mixture of two or more substances.

Ø  A solute is defined as the substance that dissolves in a solution.

Ø  A solvent is defined as the material that dissolves the other substance(s) in a solution. It is the dissolving medium.

Ø  Solubility is defined as the maximum amount of a substance (solute) which will dissolve in a given amount of solvent at a specific temperature.

            Solubility is the property of a solid, liquid, or gaseous chemical substance called solute to dissolve in a liquid or gaseous solvent to form a homogeneous solution of the solute in the solvent.

            The solubility of a substance depends on the used solvent as well as on temperature and pressure. The extent of the solubility of a substance in a specific solvent is measured as the saturation concentration where adding more solute does not increase the concentration of the solution.

            The solvent is generally a liquid, which can be a pure substance or a mixture. The extent of solubility ranges widely, from infinitely soluble such as ethanol in water, to poorly soluble, such as silver chloride in water. The term insoluble is often applied to poorly or very poorly soluble compounds.

            Under certain conditions the equilibrium solubility can be exceeded to give a so-called supersaturated solution, which is metastable.

            According to an IUPAC definition, solubility is the analytical composition of a saturated solution expressed as a proportion of a designated solute in a designated solvent. Solubility may be stated in units of concentration, molality, mole fraction, mole ratio, and other units.

Factors affecting solubility

            The solubility of one substance in another is determined by the balance of intermolecular forces between the solvent and solute, and the entropy change that accompanies the solvation. Factors such as temperature and pressure will alter this balance, thus changing the solubility.

            Solubility may also strongly depend on the presence of other species dissolved in the solvent, for example, complex-forming anions (ligands) in liquids.

Solubility will also depend on the excess or deficiency of a common ion in the solution, a phenomenon known as the common-ion effect. Solubility will depend on the ionic strength of solutions.

Temperature

            The solubility of a given solute in a given solvent depends on temperature. For many solids dissolved in liquid water, the solubility increases with temperature up to 100 °C. In liquid water at high temperatures, the solubility of ionic solutes tends to decrease due to the change of properties and structure of liquid water; the lower dielectric constant results in a less polar solvent.

            The chart shows solubility curves for some typical solid inorganic salts (temperature is in degrees Celsius). Many salts behave like barium nitrate and disodium hydrogen arsenate, and show a large increase in solubility with temperature. Some solutes (e.g. NaCl in water) exhibit solubility which is fairly independent of temperature. A few, such as cerium(III) sulfate, become less soluble in water as temperature increases. This temperature dependence is sometimes referred to as "retrograde" or "inverse" solubility.

            The solubility of organic compounds nearly always increases with temperature. The technique of recrystallization, used for purification of solids, depends on a solute's different solubilities in hot and cold solvent. A few exceptions exist, such as certain cyclodextrins.

Pressure

            For condensed phases (solids and liquids), the pressure dependence of solubility is typically weak and usually neglected in practice. Assuming an ideal solution, the dependence can be quantified as:

            The pressure dependence of solubility does occasionally have practical significance. For example, precipitation fouling of oil fields and wells by calcium sulfate (which decreases its solubility with decreasing pressure) can result in decreased productivity with time.

 Polarity

            A very polar (hydrophilic) solute such as urea is very soluble in highly polar water, less soluble in fairly polar methanol, and practically insoluble in non-polar solvents such as benzene. In contrast, a non-polar or lipophilic solute such as naphthalene is insoluble in water, fairly soluble in methanol, and highly soluble in non-polar benzene.

            The solubility is favored by entropy of mixing and depends on enthalpy of dissolution and the hydrophobic effect.

 Applications

            Solubility is of fundamental importance in a large number of scientific disciplines and practical applications, ranging from ore processing, to the use of medicines, and the transport of pollutants.

            Solubility is often said to be one of the "characteristic properties of a substance," which means that solubility is commonly used to describe the substance, to indicate a substance's polarity, to help to distinguish it from other substances, and as a guide to applications of the substance. For example, indigo is described as "insoluble in water, alcohol, or ether but soluble in chloroform, nitrobenzene, or concentrated sulfuric acid".

            Solubility of a substance is useful when separating mixtures. For example, a mixture of salt (sodium chloride) and silica may be separated by dissolving the salt in water, and filtering off the un dissolved silica.

            Another example of this is the synthesis of benzoic acid from phenyl magnesium bromide and dry ice. Benzoic acid is more soluble in an organic solvent such as dichloromethane or diethyl ether, and when shaken with this organic solvent in a separatory funnel, will preferentially dissolve in the organic layer. The other reaction products, including the magnesium bromide, will remain in the aqueous layer, clearly showing that separation based on solubility is achieved. This process, known as liquid-liquid extraction, is an important technique in synthetic chemistry.

Solubility of ionic compounds in water

            Some ionic compounds (salts) dissolve in water, which arises because of the attraction between positive and negative charges. For example, the salt's positive ions (e.g. Ag⁺) attract the partially-negative oxygens in H₂O. Likewise, the salt's negative ions (e.g. Cl⁻) attract the partially-positive hydrogens in H₂O.

AgCl_(s) Ag⁺_(aq) + Cl⁻_(aq)

However, there is a limit to how much salt can be dissolved in a given volume of water. This amount is given by the solubility product, K_sp. This value depends on the type of salt (AgCl vs. NaCl), temperature, and the common ion effect.

Solubility of organic compounds

            The principle of polarity, that like dissolves like, is the usual guide to solubility with organic systems. For example, petroleum jelly will dissolve in gasoline because both petroleum jelly and gasoline are non-polar hydrocarbons. It will not, dissolve in ethyl alcohol or water, since the polarity of these solvents is too high. Sugar will not dissolve in benzene, since sugar is too polar in comparison with benzene.

Solubility in non-aqueous solvents

            Non polar solutes are soluble in non aqueous solvents. Most available solubility values are those for solubility in water. The reference also lists some for non-aqueous solvents.

 Quantification of solubility

            Solubility is commonly expressed as a concentration, either by mass (g of solute per kg of solvent, g per dL (100mL) of solvent, molarity, molality, mole fraction or other similar descriptions of concentration.

Density

            Density is the measure of the mass per unit volume of a material (density = mass/volume). Density is a characteristic of a substance. Mass and volume vary with size but density will remain constant. Temperature will affect the density of a substance.

 Molarity

            A useful way to express exact concentrations of solutions is molarity. Molarity is defined as moles of solute per litre of solution. Molarity is symbolized by the capital letter M. It can be expressed mathematically as follows.

Molarity (M)   =          moles of solute (n)

Liters of solution (V)

            Notice that the moles of solute are divided by the liters of solution not solvent. One litre of one molar solution will consist of one mole of solute plus enough solvent to make a final volume of one litre.

Normality

            The normal concentration is another method for expressing the concentration of solutions.

            Normality (N) is defined as the number of equivalents of solute dissolved in one litre of solution.

            One equivalent of acid is the amount of acid necessary to give up one mole of hydrogen ions in a chemical reaction. One equivalent of base is the amount of base that reacts with one mole of hydrogen ions. When expressing the concentrations of bases, normality refers to the number of available hydroxyl ions. Because hydrogen and hydroxyl ions combine on a one-to-one basis, one OH- is equivalent to one H+ ion.

            ppm expresses the concentration of a solution in units of one part of solute to one million parts solvent. One ppm equals one milligram of solute per litre of solution.

INTERMOLECULAR FORCES

INTERMOLECULAR FORCES

• Inter molecular forces are the attraction between the molecules. It is a weak bond.
• Attractions exerted by one molecule on another, such as the force of attraction between water molecules in ice.
• Attractions between atoms of the noble gas elements, helium through radon.
• Attractions between molecules of one substance and molecules of another, as when two liquids are mixed, or a molecular solid such as sugar is dissolved in a liquid.
• Attraction between molecules of one substance and ions of another, as when an ionic compound dissolves in a liquid.

            Intermolecular forces (forces between two molecules) are weak compared to the intramolecular forces (forces keeping a molecule together). For example, the covalent bond present within HCl molecules is much stronger than the forces present between the neighbouring molecules. These forces exist between molecules when they are sufficiently close to each other. The forces consist of following types:

1. Dipole-dipole interactions
2. Hydrogen bonds
3. Dispersion forces
4. Vander walls forces
5. Ion–dipole forces

6. Instantaneous dipole-induced dipole forces or London dispersion forces.

Dipole–dipole interactions

            Dipole–dipole interactions are electrostatic interactions of permanent dipoles in molecules. These interactions tend to align the molecules to increase the attraction (reducing potential energy).

An example of a dipole–dipole interaction can be seen in hydrogen chloride (HCl):

            The positive end of a polar molecule will attract the negative end of the other molecule and cause them to be arranged in a specific arrangement. Polar molecules have a net attraction between them. For example HCl and chloroform (CHCl₃)

            Intermolecular forces that operate between neutral molecules having molecular dipole moments are called dipole-dipole forces. The dipole moments of two neighbouring molecules tend to align with the + end of one dipole near the - end of the other, so that forces of attraction between them are maximized.

            Such forces are obviously much weaker than those operating in ionic or covalent network solids, and give rise to potential wells having depths in the approximate range 5-20 kJ/mole. Many molecular substances with dipolar molecules exist as liquids at ambient temperature, and have relatively low boiling points. In particular, many organic compounds are of this type.

·         For Ex - SiF₄, CHCl₃, CO₂, SO₂ experience dipole-dipole intermolecular forces.

Hydrogen bonding

            Hydrogen bond is the strong dipole-dipole attractions between hydrogen atoms bonded to small, strongly electronegative atoms (O, N, or F) and nonbonding electron pairs on other electronegative atoms.

1) Bond dissociation energy of about 4-38 KJ mol–1 (0.96-9.08 Kcal mol–1).

2) H-bond is weaker than an ordinary covalent bond; much stronger than the dipole-dipole interactions.

3) Hydrogen bonding accounts for the much higher boiling point (78.5 °C) of ethanol than that of dimethyl ether (–24.9 °C).

4) A factor (in addition to polarity and hydrogen bonding) that affects the melting point of many organic compounds is the compactness and rigidity of their individual molecules.

            Adjacent molecules of the compound containing an O-H bond will be attracted to each other by virtue of these opposite charges. This force of attraction is known as the Hydrogen Bond. Usually a hydrogen bond is represented by a dotted line (Fig. 5.39).

            Intermolecular hydrogen bonding is responsible for the high boiling point of water (100°C) compared to the other hydrides that have no hydrogen bonds. Intra molecular hydrogen bonding is partly responsible for the secondary, tertiary, and quaternary structures of proteins and nucleic acids. It also plays an important role in the structure of polymers, both synthetic and natural. Hydrogen bond forces cause potential wells of depth in the range 5-50 kJ/mole.

 Examples  -  CHCl₃, CH₃CH₂OH, HNO₃, PH₃

            It is understandable that substances having nearly the same molecular weights, have the same boiling point. The boiling points of alkanes and ethers of comparable molecular weights are not far apart, but the boiling points of alcohols having almost equal molecular weights are considerably higher.

            This can be explained on the basis of hydrogen bonding. Ethanol forms hydrogen bonds. Extra energy in the form of heat is required to break the hydrogen bonds holding the molecules together before it can be volatilized. Propane and dimethyl ether do not form hydrogen bonds and, therefore, have low boiling points.

Effect on Water-Solubility. A hydrogen-bonded substance is usually soluble in another hydrogen bonded substance.

            For example, alcohols are soluble in water but alkanes are not. This is because a non polar alkane molecule cannot break into the hydrogen bonded sequence in water . It Cannot replace the hydrogen bonds that would have to be broken to let it in.

            An alcohol molecule is capable of hydrogen bonding. It can slip into the hydrogen bonded sequence in water. It can replace the hydrogen bonds that must be broken to let it in.

            Thus alcohols of low molecular weight are water soluble. However, when the alkyl group is four or more carbons in length the alkane nature of the molecule predominates, and water solubility fans off sharply. Alcohols containing more than seven carbons are insoluble in water.

London dispersion forces

            Otherwise known as quantum-induced instantaneous polarization or instantaneous dipole-induced dipole forces, the London dispersion force is caused by correlated movements of the electrons in interacting molecules. The electrons, which belong to different molecules, start "feeling" and avoiding each other at the short intermolecular distances, which is frequently described as formation of "instantaneous dipoles" that attract each other.

Debye (induced dipole) force

            The induced dipole forces appear from the induction (also known as polarization), which is the attractive interaction between a permanent multipole on one molecule with an induced (by the former di/multi-pole) multipole on another. This interaction is called Debye force after Peter J.W. Debye.

            The example of an induction-interaction between permanent dipole and induced dipole is HCl and Ar. In this system, Ar experiences a dipole as its electrons are attracted (to H side) or repelled (from Cl side) by HCl. This kind of interaction can be expected between any polar molecule and non-polar/symmetrical molecule. The induction-interaction force is far weaker than dipole-dipole interaction, however stronger than London force.

Van der Waals Forces

            In addition to chemical bonding between atoms, there is another type of attractive force that exists between atoms, ions, or molecules known as van der Waals forces.

            These forces occur between the molecules of non polar covalent substances such as H₂, Cl₂ and He. These forces are generally to be caused by a temporary dipole, or unequal charge distribution, as electrons constantly move about in an atom, ion, or molecule. At a given instant, more electrons may be in one region than in another region, as illustrated in Figure.

            The temporary dipole induces a similar temporary dipole on a nearby atom, ion, or molecule. Every instant, billions of these temporary dipoles form, break apart, and reform to act as a weak electrostatic force of attraction known as van der Waals forces.

            It is important to note that van der Waals forces exist between all kinds of molecules. Some molecules may have these forces, as well as dipole or other intermolecular forces. Therefore, the strength of the van der Waals forces between substances increases with increasing gram molecular mass.

INTRA MOLECULAR FORCES

INTRA MOLECULAR FORCES

• Intra molecular forces are the attraction of atoms within the molecule. It is a strong bond.

·  An intra molecular force is any force that holds together the atoms making up a molecule or compound.

• Attraction between positive and negative ions in a crystal of an ionic compound.
• Covalent bonds in molecular substances.
• Metallic bonds

Types of Bonds

There are four basic ways in which chemical combination occurs:

(1) Ionic or electrovalent bond

(2) Covalent bond

(3) Coordinate bond

(4) Metallic bond

Ionic or Electrovalent Bond.

            Ionic or electrovalent bonds are formed by transfer of valence electrons from one atom to another. This type of bond unites two atoms one of which has excess electrons than the stable number (2 or 8), and the other is short of electrons.

For Ex - Sodium chloride is a typical compound formed in this way.

            Here the sodium atom (2, 8, 1) transfers its excess electron to chlorine atom (2, 8, 7), and thus both attain a stable inert gas type electron configuration. Sodium acquires the electron configuration of neon (2, 8) and becomes positively charged. Chlorine acquires the electron configuration of argon (2, 8, 8) and becomes negatively charged. These oppositely charged ions are held together by electrostatic force of attraction. This type of bond is commonly found in inorganic compounds.

            Ionic or electrovalent compounds are non volatile, soluble in water and possess high melting points. Their aqueous solutions conduct electric current.

            The sodium atom loses the one electron in its outer shell to the chlorine atom, which uses the electron to fill its outer shell. When this occurs, the sodium atom is left with a +1 charge and the chlorine atom a -1 charge. The ionic bond is formed as a result of the attraction of the two oppositely-charged particles.

Covalent Bond.

Ø  Covalent bonds are formed by mutual sharing of electrons. This type of bond unites two atoms, both of which are short of electrons. The two atoms contribute one electron each and then share the resulting pair of electrons.

For Ex -  Hydrogen is the simplest compound formed in this way.

       H. + xH -----------H ~H or H-H

Here the two electrons are shared and give to each hydrogen atom the configuration of helium.

Ø  This type of bond is termed covalent bond and is indicated by a line. Covalent bonds are commonly found in organic compounds.

Ø  A covalent bond between two atoms results from the overlap of an orbital of one atom with an orbital of another atom. When two orbitals overlap they share the same region in space and a new orbital, called a Molecular Orbital (MO), is formed.

Ø  Covalent compounds are volatile, generally insoluble in water, but soluble in organic solvents. They possess low melting and boiling points. Their solutions do not conduct electric current.

Ø  A covalent bond is formed only if: (1) The two combining orbitals are half-filled and      (2) The bonding orbitals approach each other in proper alignment needed for an effective overlap. The electrons in the bonding orbitals have opposite spins.

Types Of Covalent Bonds

There are two main types of covalent bonds or molecular orbitals:

(1) Sigma (α)bonds

(2) Pi(π) bonds

Sigma Bond

A sigma bond is formed by the linear or end-to-end overlap of orbitals. They may be

Obtained

a) By the overlap of two s orbitals,

b) By the overlap of Px and an s orbital,

c) By the overlap of two p orbitals,

 Sigma bonds are symmetrical around the line drawn between the two nuclei. They are represented by a single line drawn between the atomic symbols. The electrons that occupy a sigma bond are called Sigma (α) Electrons.

 Pi-Bond.

A pi (π) bond is formed by parallel or side-to-side overlap of orbitals. For example,

             Like the p orbitals from which it is obtained, a π bond has two lobes. One half of the π bond lies above the plane containing the two nuclei and the other half lies below. The electrons that occupy π bond are known as the Pi (π) Electrons.

            There are two important differences between sigma and pi bonds.

(1) Pi electrons are loosely held than a pair of electrons in a α bond. They are of higher energy. As a result, π bonds are more easily broken and are more reactive than α bonds.

(2) Rotation of atoms is not possible around a π bond. If any attempt is made to do so, the lobes of p orbitals will no longer be coplanar and will not overlap to form the π bond. This restriction in rotation around a π bond is responsible for Cis and Trans isomerism in alkenes.

            Unlike an ionic bond, a covalent bond holds together specific atoms. Covalent bonding can be single covalent, double covalent, or triple covalent depending on the number of pairs of electrons shared. Figure 7 shows the bonding that occurs in the methane molecule, which consists of four single covalent bonds between one carbon atom and four hydrogen atoms.

Carbon-Carbon Single Bond

             Carbon atom has the wonderful property of uniting with other carbon atoms through covalent bonds. This serves to construct the carbon structure of organic molecules.

For example - The molecules of hydrocarbons, ethane and propane contain two and three carbon atoms respectively linked by covalent bonds.

Carbon-Carbon Double Bond

             In some compounds, two of the valencies of a carbon atom may be satisfied by union with the two valencies of another carbon atom.

For example - In ethylene the two carbons are united by two covalent bonds.

Carbon-Carbon Triple Bond

             Sometimes two adjacent carbons are linked together by three covalent bonds.

For example - acetylene

Coordinate Bond.

            Coordinate bond is also formed by mutual sharing of electrons but in this case the two electrons that are shared come from the same atom. A coordinate bond unites two atoms, one of which has a spare pair of electrons and the other is short of a pair of electrons. The first atom (donor atom) contributes one pair (lone pair) of electrons and the second atom (acceptor atom) accepts it.

            After the formation of the bond, the lone pair of electrons is held in common. The coordinate bond is represented by an arrow, pointing away from the donor atom.

For example - the coordinate bond is found in the boron hydride-ammonia complex.

            When both shared electrons in a covalent bond come from the same atom, the bond is called a coordinate covalent bond. Coordinate covalent bond is a single bond similar in properties to a covalent bond.

 For example- negatively-charged chlorate ion.

            The ion consists of one chlorine atom and three oxygen atoms with a net charge of -1, and is formed with two coordinate covalent bonds and one covalent bond. The chlorine atom has effectively gained an electron through the covalent bond, which causes the overall negative charge.

Metallic Bond

                 Metallic bonds are bonds where the atoms achieve a more stable configuration by sharing the electrons in their outer shell with many other atoms.

            Elements of similar but relatively low electro negativities (i.e., metals) form metallic bonds. It is these bonds that are operative in a typical metal and that are responsible for metallic properties: reflectivity, conductivity, malleability, and strength. Metallic bonds form between atoms that have fewer valence electrons than they have valence orbitals.

            It would contribute its valence electron to a bond with one or the other of the pair. The sodium atoms are therefore driven to aggregate, using their relatively few valence electrons to bond together as many nuclei as possible via the Coulomb attraction. The aggregation is very regular in nature, resulting in an ordered arrangement of metal atoms called a lattice. As a result of aggregation, each sodium atom uses all of its valence orbitals in forming partial bonds with a number of neighbours.

            The nuclei are represented as points, and the electrons as waves. That the electrons are delocalized waves is consistent with the conductivity of metals, and with their flexibility and strength. Putting stress on metals moves the nuclei, but the waves adjust, maintaining bonding. Because the electrons are waves, they are everywhere in the crystal at once. Thus the electrons absorb and reemit photons of all energies in the visible region of the spectrum, giving metals their highly reflective appearance.

Polar and non polar Bonds

            When a covalent bond is formed between two atoms with different electro negativities, the electrons involved in the bond are not shared. The atom with higher electro negativity pulls the bonding electrons closer to it.

            In other words, the electron density in the molecular orbital would be greater around the atom with higher electro negativity. The result of this displacement of the molecular orbital toward the more electronegative atom will be that the more electronegative atom will acquire a small negative and the less electronegative atom will acquire a small positive charge. This is called polar bond.

                        Let us now consider the C-CI covalent bond. Because chlorine is more electronegative than carbon, the electro n density in the molecular orbital would be higher around chlorine atom than around the carbon atom. Thus the chlorine atom acquires a small negative charge and the carbon atom acquires a small positive charge. This may be indicated in the following way.

            When a covalent bond is formed between two atoms with same e1ectronegativities (C-C, H-H, and F-F), the electrons involved in the bond are shared equally. Electron density in the molecular orbital binding the two atoms together is same around each atom. There is no positive and negative end. Such a bond is said to be a Non polar Bond.

Polar and Non polar Molecules

            When the shared pair of electrons which are forming the bond in a molecule are not shared equally, the resulting molecule will have a positive end and a negative end. This type of bond is a polar covalent bond. The molecules are called dipolar or polar molecules.

            A molecule is said to be a polar molecule if it satisfies the following two conditions:

(1) The molecule must contain one or more polar bonds.

(2) The polar bonds must be so directed that there are separate centres of positive and negative charges in the molecule.

For example –

a)      Carbon dioxide (CO₂) has two polar carbon-oxygen bonds, but the molecule is not polar. The carbon dioxide molecule is linear and the centre of the two negative charges is at the same place as the centre of the positive charge - the carbon atom (Fig. 5.35).

 

b)      Water molecule has two polar oxygen-hydrogen bonds. The H-O--H bond angle is 104°. The centre of the positive charge is between the hydrogen atoms (Fig. 5.36). As the centre of positive charge and the centre of negative charge do not coincide, water molecule is polar.

c)      Carbon tetrachloride has four polar carbon-chlorine bonds, but the molecule is not polar. The Carbon tetra chloride is tetrahedral (all Cl-C-Cl angles are 109°28'), and the centre of four negative charges is at the same place as the centre of the positive charge of the carbon atom (Fig. 5.37).

Examples

Dipole moment 

            Polar molecules behave like small magnets and possess a dipole moment. The dipole moment (µ) is the product of the magnitude of the charges and the distance between the charges.

            The measure of molecular polarity is a quantity called the dipole moment (µ). Dipole moment is defined as magnitude of charge (Q) times distance (r) between the charges.

µ = (Q)(r)       Q charge in coulombs (C)      r distance in meters (m)

            The product of the strength of either of the charges in anelectric dipole and the distance separating the two charges. Itis expressed in coulomb meters. Dipole moment is a vectorquantity;  its direction is defined as toward the positive charge.

            Many molecules have such dipole moments due to non-uniform distributions of positive and negative charges on the various atoms. Such is the case with polar compounds like hydrogen fluoride (HF), where electron density is shared unequally between atoms.

            A molecule with a permanent dipole moment is called a polar molecule. A molecule is polarized when it carries an induced dipole. The physical chemist Peter J. W. Debye was the first scientist to study molecular dipoles extensively, and, as a consequence, dipole moments are measured in units named debye in his honor.

With respect to molecules, there are three types of dipoles:

Permanent dipoles: These occur when two atoms in a molecule have substantially different electro negativity: One atom attracts electrons more than another, becoming more negative, while the other atom becomes more positive.

Instantaneous dipoles: These occur due to chance when electrons happen to be more concentrated in one place than another in a molecule, creating a temporary dipole.

Induced dipoles: These can occur when one molecule with a permanent dipole repels another molecule's electrons, "inducing" a dipole moment in that molecule.

            The simplest example of a dipole is a water molecule. A molecule of water is polar because of the unequal sharing of its electrons in a “bent” structure. Since oxygen has a higher electro negativity than hydrogen, the side of the molecule with the oxygen atom has a partial negative charge while the hydrogen, in the center, has a partial positive charge. Because of this, the direction of the dipole moment points towards the oxygen.

It is expressed in the units of Debye and written as D (where 1 Debye = 1 x 10^-18e.s.u cm). nA dipole moment is a vector quantity and is therefore represented by a small arrow with a tail at the positive center and head pointing towards a negative center. In the case of a Water molecule, the Dipole moment is 1.85 D, whereas a molecule of hydrochloric acid is 1.03 D and can be represented as: