There are two types of forces present in molecules:
1. Intra-molecular Forces.
2. Inter-molecular Forces.
Intra-molecular Forces Vs Inter-molecular Forces.
Summary of Intermolecular Forces or van der Waal’s Forces
1. Intramolecular Forces
Definition
Intramolecular forces hold atoms together in a molecule i.e. they operate within the molecules or fundamental units of substances.
Other Name
They are also called chemical bonds or primary bonds which are of two types namely ionic and covalent bonds. (e.g. water molecule consists of 2 hydrogen and 1 oxygen atom joined together through shared electron pair called covalent bond while sodium chloride (formula unit) consists of 1 Na+ and 1 Cl– ions held together by electrostatic forces of attraction known as Ionic Bond).
Importance
Intramolecular forces stabilize individual molecules.
Strength
intramolecular forces are almost 40-400% times the strength of intermolecular forces.
2. Intermolecular Forces
Definition
Intermolecular forces or IMFs or secondary forces or more famously as van der Waal’s forces are the attractive forces between the neutral molecules holding them together at certain temperature. As the name implies, they operate or exist between molecules of covalent compounds (not within molecules, bonds exist within molecules) and the atoms of monoatomic elements i.e. noble gases. (Some elements such as noble gases exist with IMFs and no bonding at all.) These attractive forces are actually physical bonds and are electrostatic in nature.
Other Name
Intermolecular forces are also known as van der Waal’s Forces named after the Dutch Physicist (1837-1923) Johannes van der Waal who first described them.
Importance
IMFs play an important role in determining the properties of substances.
Intermolecular forces account for why a substance has different states at different temperature (i.e. condensed state of matter), non-ideal behaviour of gases and liquefaction of gases. Stronger IMFs makes the molecules closer together forming liquids and even stronger IMFs pack the molecule together forming solids. Thus the three states of matter represent three different relative magnitudes of IMFs. Many physical properties are related to the strength of IMFs and their values thus manifest the strength of IMFs e.g. state, solubility, evaporation, vapour pressure, latent heats, specific heats, fluidity, viscosity, surface tension, boiling points, melting points, freezing point, dipole moment etc.
Occurrence
IMFs are only associated with covalent compounds and they are not encountered in ionic compounds.
Nature of Forces
All IMFs are electrostatic (electrical) forces involving mutual attraction between unlike charges i.e. positive and negative charges (between an area of negative charge on one molecule and an area of positive charge on a second molecule) or the mutual repulsion of like charges. Polar molecules have partial positive and partial negative charges on their ends (dipoles). These dipoles strongly influence the intermolecular forces. [The greater the polarization of the bond, stronger will be the attraction between the molecules i.e. IMFs.]
Strength
Generally, IMFs are much weaker than intra-molecular forces or primary bonds (i.e. ionic or covalent bonds). Their bond strength range from 1-40 kJ/mol (1-10 or 1-16 kcal/mol). Generally, intermolecular forces tend to be less than 15% as strong as covalent or ionic bond.
Types of IMFs
There are three types of IMFs:
1. London Forces (exist in all substances especially in non-polar compounds).
2. Dipole-dipole forces (exist in all polar compounds lacking Hd+)
3. Hydrogen bonding (exist in all polar compounds containing Hd+)
4. Ion dipole interaction (exist in solutions of ionic compounds)
Definition
These forces were first reported by German physicist Fritz London. All particles, whether individual atoms, molecules or ions experience dispersion forces which result from the motion of electrons around atom. Dispersion forces are also called London Forces in the honour of German physicist Fritz London who first recognized and identified them in 1930.They are also known as van der Waal’s forces. These attractive forces exist among non-polar molecules which becomes polar temporarily.
London dispersion forces are a temporary attraction between two adjacent atoms. One atom's electrons are unsymmetrical, which creates a temporary dipole. This dipole causes an induced dipole in the other atom, which leads to the attraction between the two.
Examples
gasoline and benzene, are non-polar molecules but exist in liquid state at room temperature due to the presence of London dispersion forces.
Occurrence
Dispersion forces operate between all molecules whether they are polar or non-polar. However, the dispersion forces are the only attractive forces in non-polar molecules (F2, Cl2, Br2, I2, H2, O2 etc.) and atoms of noble gases (like He, Ne, Ar, Kr, Xe).
Dispersion forces are strong when particles are close together but rapidly weakens as they move apart i.e. dispersion forces are effective only at very short distances. Thus, dispersion forces are significant only at low temperature.
Bond Strength
Dispersion forces are the weakest of all intermolecular forces. Their bond strength (energy) is of the order of 1-5 kJ/mol (1-10 kJ/mol).
Mechanism of Origin of London Forces
These attractive forces exist among non-polar molecules which becomes polar temporarily. Nevertheless, gasoline and benzene, are non-polar molecules but exist in liquid state at room temperature due to the presence of London dispersion forces.
When non-polar substances such as H2, Cl2, F2, CH4, He, Ne, Ar etc. are allowed to liquefy by lowering the temperature, a temporary interaction is developed among their molecules due to the distortion of electrons cloud of one atom by the electronic influence of other atom. This makes a short-lived polarization of molecule and it is said to be instantaneous dipole. This instantaneous dipole then distort the electron density of nearer atom and hence produce an induce dipole in the nearby atom.
Details
Dispersion forces result from attraction between temporary dipole. The distribution of electrons around the nucleus of an atom (like neon or non-polar molecule) is spherically symmetrical (uniform). However (especially when atoms or molecules come extremely close together), the vibration of electron clouds at any given instant may be more concentrated at one side of the atom producing a small transient instantaneous dipole (although net value of these dipoles is zero because instantaneous dipoles cancel one another). These tiny temporary dipoles on one atom (or in a molecule) can induce a similar dipole called induced dipole on an adjacent atom as electrons of one atom repel electrons of neighbouring atom. This polarizes each atom (or molecule) which are oriented in such a way that their oppositely charged ends (regions) face each other thereby causing the atoms to be attracted to each other giving rise to weak attractive forces called Dispersion Forces. [Since dispersion forces exist between Instantaneous dipole and induced dipoles, they are also referred to as Instantaneous dipole-Induced dipole Interaction.]
Factors Affecting London Forces
The strength of London dispersion forces depends mainly upon following factors:
a). Molecular Wt./Molecular size (direct relation; (LDF tend to increase with increasing molar mass)
b). Number of atoms in molecules
c). Molecular Shape
d) Polarizability
(a) Size of atom or molecules or Atomic or Molecular size
The strength of dispersion forces depends upon the size of the electronic cloud of the atom or molecule. Large size atom or molecules have greater number of electrons therefore more distortion of electrons is possible i.e. the dispersion becomes easy which increases the strength of London forces. Greater the size of molecules, stronger the London forces between them, because the dispersion of electrons becomes easier.
For example
London forces increase down the group in halogens and in inert gases due to their increasing size. Melting and boiling points also increase down the group due to same reason.
The elements of the zero group in the periodic table are atomic gases. They don’t make covalent bonds with other atoms because their outermost shells are complete. Their boiling points increase down the group from helium to radon due to increasing London forces owing to their increasing atomic size.
(b) Number of atoms in molecules
Another important factor that affects the strength of London forces is the number of atoms in a non-polar molecule. The greater the number of atoms in a non-polar molecule, the more is the number of electrons and greater is the polarizability of molecules hence the stronger the electronic distortion which results in enhancing the strength of London forces.
For example
the intermolecular attractive forces are stronger in H2 than in He due to greater number of atoms in H2.
(c) Shape of the molecules
Shape also plays a role in the magnitude of LDF. The strength of LDF increases with increasing molecular surface. The larger surface increases the chances for the induced charge separation.
The linear molecules having greater surface area than non-linear molecules would have stronger LDF. Thus linear longer, less compact molecules have stronger LDF than more compact molecules.
For example
n-pentane with longer less compact molecular structure has stronger LDF (manifested by high boiling point of 309.4 K) than neo-pentane with more compact molecular structure (manifested by low boiling point of 282.7 K). Thus LDF decreases with molecular branching and increases with molecular symmetry.
Long straight chain molecules can come very close to each other than branched chain molecules. Therefore, London forces are stronger in straight chain molecules.
For example
Therefore, London forces are stronger in n-butane than in iso-butane.
(d) Polarizability
LDF depends on polarizability of molecules. polarizability is the quantitative measurement of the electronic distortion and it is the extent to which the electronic cloud can be polarized. Polarizability is the ease with which an external nearby electric field can distort a molecule’s electron cloud and can induce a dipole (by altering the electron distribution) within a molecule. The greater the polarizability of a molecule, the more easily its electron cloud can be distorted to induce a momentary dipole and the stronger the LDF.
A smaller molecule or a lighter atom is less polarizable and has smaller LDF because it has only a few tightly held electrons.
A large molecule or heavier atoms tends to have greater polarizabilities and LDF because it has many electrons (so there is a higher probability of asymmetric electron distribution) which are less tightly held and are farther from the nucleus (so many asymmetric distribution produces a larger dipole due to larger charge separation).
Examples
1. The increasing polarizabilities of noble gases from upper to the downward direction are responsible for the increasing melting and boiling points.
As the atomic number increases down the group the outermost electrons become away from the nuclei. The dispersion of the electronic clouds becomes more and more easy. So the polarizability of these atoms goes on increasing.
2. Similarly, the boiling points of halogens in group VIIA go on increasing from fluorine to iodine. All the halogens are non-polar diatomic molecules, but there is a big difference in their physical states at room temperature. Fluorine is a gas and boils at188.1°C. The polarizability of iodine molecule is much greater than that of fluorine.
3. The increasing polarizabilities for alkanes from lower molar masses to higher ones are due to the big sizes of molecules form methane onward and is responsible for the increasing m.p. and b.p. downward.
LDF a Size a Polarizability a Atomicity
Definition
Liquid molecules are either polar or non-polar. Polar molecules bear partial positive and partial negative charges due to electronegativity differences on their opposite ends (dipoles). Dipole-dipole forces exist between neutral polar unsymmetrical molecules that have permanent net dipole moments (i.e. that have permanent separation of positive and negative charge).
These forces are created when the positive end of one molecule is attracted by the negative end of other molecule. Therefore, these forces are electrostatic in nature.
Examples
Dipole-dipole interactions (DDI) are significant only when molecules are in close contact. DDI exists in SO2, SCl2, PCl3, CH3Cl, HCHO, acetone in acetone, triethyl amine in acetone, etc. They exist in all three phases i.e. solid, liquid and gases. Dipole-dipole forces are more operative in liquids than gases.
An example of liquid containing dipole-dipole forces is iodine mono chloride (reddish brown oily liquid). The more electronegative chlorine atoms bears partial negative charge while less electronegative iodine atom acquires partial positive charge. These opposite poles of molecules attract each other through electrostatic attraction creating dipole-dipole interaction making the compound liquid.
HCl is a polar molecule and there is a partial positive charge on hydrogen atom at one end and partial negative charge on chlorine atom at the other end. These opposite poles attract each other in the liquid states.
Bond Strength
DDF are generally weaker than hydrogen bond (and ion-dipole forces) but stronger than dispersion forces. Their bond strength (energy) is in the range of 3-10 kJ/mol (0.1-10 kJ/mol). Dipole-dipole forces are only 1% as effective as covalent bonds i.e. as a rough comparison, their bond strength is about 1% of a normal covalent bond.
Factors Affecting DDF
1. Polarity or Difference in EN or dipole moment
2. Intermolecular distance
The strength of dipole-dipole forces depends on the dipole moment, which in turn depends on polarity, which is measured by the difference of electronegativities of polarized bonded atoms in the polar molecules. The greater the difference in electronegativities, the stronger the bond polarization (polarity) and greater the dipole-dipole forces.
These forces are affected by the distance between the molecules, therefore these forces are weaker in gases.
Mechanism of Origin of Dipole-Dipole Force
The hydrogen chloride (HCl) molecule is polar due to greater electronegativity of chlorine than hydrogen. Thus, H–Cl bond is polarized creating a partial positive charge on hydrogen atom and a partial negative charge on chlorine atom thereby producing permanent dipole moment (Hδ+–Clδ–). On approaching each other, an electrostatic force is created between polar H–Cl molecules due to orientation of δ+ end of one HCl molecule towards the δ– end of neighbouring molecule that gives rise to an attractive force in addition to dispersion forces which causes dipolar HCl molecules to be attracted to each other giving rise to dipole-dipole forces.
The presence of DDF is manifested by the rise in boiling point of HCl. Due to presence DDF polar HCl boils at –85°C but non-polar F2 boils at –188°C though both have nearly same molecular weights (molecular weight of HCl is 36.5 amu and that for F2 is 38 a.m.u.) so both have same strength of dispersion forces.
Effects of DDF
The presence of these forces increases the melting and boiling point of the substances. That is why polar substances have high melting and boiling points than non-polar substances.
Definition
Besides normal or primary bonds (such as ionic and covalent), there exist some secondary bonds known as Hydrogen Bond.
Hydrogen bond is an unexceptionally (unusually) strong and special type of dipole-dipole interaction. Hydrogen bondings differ from the word ‘bond’ since it is an intermolecular force or secondary bond not an intramolecular force or primary bond as in the common use of word bond.
Hydrogen bonding is also known as Protonic Bridge as hydrogen or protons (Hδ+) causes the association of molecules in larger aggregates acting as a bridge between polar molecules. It is expressed with dotted lines (………………. ) like this —X—H…Y—
Mechanism of Origin of Hydrogen Bonding/ Examples with explanation
In certain polar molecules like HF, H2O, NH3 etc. where the H atom is bound with highly electronegative smaller size atoms non-metallic atoms like, F, O or N, the bond between H–F, O–H and N–H is polarized creating a partial positive charge on H atom and a partial negative charge on the electronegative atom (F, O, N).
When such molecules come close to each other, an electrostatic attraction is set up between the dipole of such polar molecules in the form of unusually strong and special type of dipole-dipole interaction which binds electropositive H atom (+ive pole) of one molecule to electronegative atom (–ive pole) of other molecules. This type of strong attractive force involving hydrogen and causes the association of molecules in larger clusters as (HF)x, (H2O)x, (NH3)x is known as Hydrogen Bond or Protonic Bridge. It is expressed with dotted lines (……………….).
Factors affecting the strength of hydrogen bond
The strength of hydrogen bond depends upon the polarity or ionic character of molecules, which in turn depends on the difference in electronegativities. Greater the difference in electronegativities, stronger will be the strength of hydrogen bonding. That is why the strongest hydrogen bond exists in HF (as F is the most electronegative atom). The second strongest hydrogen bonding occurs in water.
The strength of hydrogen bonding also depends on the size of the electronegative atom e.g. larger Cl and S atoms whose electronegativities are almost the same as that of N atom, form hydrogen bond to a lesser degree.
Bond Strength/ Strength of Hydrogen Bond
Hydrogen bond is the strongest among all intermolecular forces however it is much weaker than ordinary covalent or ionic bond (whose bond energy is about 150-500 kJ/mol). The strength of hydrogen bond is about 5 to 10% of the strength of covalent bond (i.e. it is only about 1/10 to 1/15 as strong as covalent bond i.e. it is almost 10 times weaker than covalent bond).
Hydrogen bond is the strongest among all intermolecular forces (but less stronger than IDI). The bond strength (energy) of hydrogen bond is of the order of 10-40 kJ/mol (3-10 kcal/mole).
Bond energy of H -------- F hydrogen bond = 41.8 kJ/mol
Bond energy of H -------- O hydrogen bond = 29.4 kJ/mol
Bond energy of H -------- N hydrogen bond = 8.4 kJ/mol
Reason of High Strength of Hydrogen bond
The high strength of hydrogen bond is due to lack of inner core electrons to shield its nucleus producing a virtually naked nucleus and also electron-poor (or electron-deficient) hydrogen atom has a small size, so it can approach an electronegative atom very closely and thus interact strongly with it.
Hydrogen bonding is stronger IMF of attraction due to following reason:
(i) Sufficient positive charge on H atom due to almost complete shifting of the shared pair of electrons towards the more electronegative atom in polar molecule
(ii) Absence of inner non-bonding electron to create force of repulsion giving close approach of H to other atom.
In some case, atoms other than F, O or N can also form Hydrogen donning like in chloroform (CHCl3). The three Cl atoms take away the electrons of carbon which creates electron deficiency on C atom which make H atom electron deficient. Thus, this H atom of chloroform form strong H-bond with oxygen atom of the acetone.
Bond Lengths of Hydrogen bonds
Bond lengths of hydrogen bond depends upon the nature of molecules. A normal covalent bond has an average bond length of 0.96 Å while bond lengths of hydrogen bonds is of the order of 1.97 Å .
In water, O–H -------- F bond distance = 2.76Å
In water, F–H -------- F bond distance = 1.57 Å (2.55 Å)
Types of Hydrogen Bond
1. Intermolecular Hydrogen Bond
2. Intramolecular hydrogen bond
1. Intermolecular Hydrogen Bond
It takes place between two or more than two same or different molecules.
For example, in water, ammonia, HF, H2O2, alcohols, carboxylic acids, acetic acid, amines etc.
2. Intramolecular hydrogen bond
It takes place between the atoms of the same molecule i.e. within the molecule (intra means within). [This type of hydrogen bonding may lead to the linking of two groups to form six-membered ring structure called Chelation. Such type of hydrogen bonding is extensively seen in ortho isomers of aromatic organic compounds (the occurrence of intramolecular hydrogen bonding is not possible in meta and para isomers of aromatic organic compounds because of the large size of the ring or chelation].
For example in orthonitrophenol, orthohydroxybenzaldehyde, orthohydroxybenzoic acid (salicylic acid), orthonitrobenzoic acid, maleic acid etc.
Due to presence of intramolecular hydrogen bonding, orthonitrophenol has different properties from those of its other isomers (i.e. meta and para isomers).
e.g.
i. Melting of orthonitrophenol is 214°C while those of its m-isomer & p-isomer are 290°C & 279°C respectively.
ii. orthonitrophenol is volatile in steam and less soluble in water than the other two isomers.
Importance of hydrogen Bonding/Applications of Hydrogen bonding
Hydrogen bonds are very important for life on earth and have a wide range of applications in many esseintal chemcial and biochemcial process.
For example:
(i) Hydrogen bonding plays a major role in biological (life) processes due to its weak nature. Water is an essential need of human beings and other animals. It exists in liquid state at ordinary temperature rather than gas due to the presence of hydrogen bonding. Contrarily hydrogen sulphide is a hydride of same group as water but exist in gaseous state at room temperature due to lack of hydrogen bond.
(ii)Macro biomolecules like proteins, deoxyribonucleic acid (DNA) etc. play a vital function for life. Hydrogen bond hold their chain in particular sequence.
(iii) Cleaning actin of soap and detergent is based on hydrogen bond formation between the polar part of soap and detergent with water molecule.
(iv). The adhesive action of paints on surface is developed due to hydrogen bonding.
(v) Fabric of cloths is made of fibers such silk, polyester, nylon etc. The rigidity and tensile effect in threads is devolved by hydrogen bonding.
Explanation of Properties of Water in terms of Hydrogen bonding/Consequences of H-Bonding (Effects of H-Bonding on Physical Properties)
Some unique properties of water are attributed due to the presence of hydrogen bonding among its molecules.
High Specific heat
Hydrogen bonding greatly increases the value of specific heats of compounds e.g. water has high specific heat of 4.2 J/g due to presence of hydrogen bonging.
The amount of heat energy required to raise the temperature of one gram of any substance by 1oC is called specific heat. Water has high specific heat due to hydrogen bonding due to which water does not warm or cool rapidly.
High specific heat of water has marked effect on the weather. It plays important role in moderating the temperature of earth’s surface. The vast of amount of water on the surface of the earth, thus acts as a Giant Thermostat to moderate the earth’s temperature. Due to its high specific heat, water is neither cooled quickly nor is heated quickly i.e. water undergoes temperature changes very slowly. Heat is absorbed and stored at day times while sun shines but released at nights. This is why the temperature of coastal areas usually remains moderate throughout day i.e. the coastal areas and the areas near large lakes have temperate climate due to slow rise or fall of temperature of water during the day or at night.
But in desert areas temperature usually shoots up at day time and fall steeply at nights because rocks and sands have lower specific heat.
Its high specific heat explains its use in hot water bottles and in radiators of heating rooms.
High Boiling Point of water/ Effect on Melting and Boiling Points/Volatility
Compounds containing HB are less volatile than the substances having no HB. Hydrogen bonding greatly elevates melting and boiling points of compounds. e.g. water has a melting point at least 100°C higher than expected on the basis of the melting point of hydrides of group VI elements like H2S, H2Se and H2Te. Similarly, water has a boiling point almost 200°C higher than expected (the boiling point of water is supposed to be 150 K or −123°C) from the boiling points of H2S, H2Se and H2Te. The elevation in melting and boiling points of water is due to strong hydrogen bonding in water which produces an electrostatic attraction b/w neighbouring water molecules. Thus more energy is required to separate the molecules from each other.
Similarly, first hydrides of group V, VI, VIIA of the periodic table (NH3, H2O, HF) have higher melting points and boiling points than other hydrides of same group (such as PH3, H2S, HCl) is present in which either no or weaker hydrogen bond is present.
The boiling point of water is quite higher as compared to other polar liquids like HF due to extensive three-dimensional hydrogen bonding. Each water molecules has two polar polarized hydrogen atoms along with one oxygen atom with tow lone pairs of electrons that allows a water molecule to form maximum four hydrogen bonds in three-dimensional space. The breaking of these extra hydrogen bonds extra energy is required that results water boils at high temperature (100oC). In HF although strength of hydrogen bonding among its molecules is stronger due to high electronegativity of fluorine atom but boiling point of HF is quite lower than water as fluorine can make less number hydrogen bonds (in fact two only) and hence in HF net attractive forces are less resulting in its low boiling point.
High Density of Water
Due to the presence of hydrogen bonds water molecules are strongly attracted and occupy less volume. Since density is inversely proportional to volume, the density of water relatively high. At 20oC, the density of water is 1 g/cm3. Marine life survives under water due to high density of water.
It is generally observed that most solids and liquids expand on heating and contracts on cooling but the expansion of water is slightly different from the other liquids. On cooling water from its boiling point (100°C), its volume goes on decreasing up to 4°C but on cooling below 4°C, its volume again starts increasing till its freezing point (0°C). Since volume of water at 4°C is minimum, therefore, its density is maximum at 4°C (1 gm/cm3). Thus on cooling water below 4°C or heating above 4°C, its volume increases and it becomes lighter and lighter. That is why the density of ice is less than that of liquid water and accounts for why ice floats on water. Thus the expansion of water is irregular and unique and is called Anomalous Expansion of water.
High viscosity of water
The viscosity of water is higher than other liquids of comparative size due to strong hydrogen bonding in water molecules. Water has multiple hydrogen bonds in three-dimensional space that results in closeness of its molecules that leads to high resistance in its flow resulting in its high viscosity.
High surface tension of water
Water has high surface tension due to the strong cohesive forces in the form extensive three-dimensional hydrogen bonding.
Anomalous Behaviour of Water
Most liquids expand upon heating and contract upon cooling. Water, between a specific temperature range, seems to break this notion. “The Anomalous Expansion of Water is the increase in the volume of water when cooled from 4°C to 0°C (or 39.2°F to 32°F ). Water’s anomalous expansion is an abnormal property in which it expands instead of contracting as the temperature rises over 4°C to 0°C, making it less dense. The density is highest at 4°C and decreases below that temperature. Because water molecules generally form open crystal formations when solidified, the density decreases as it freezes.
Explanation
The density of water increases from 0°C to maximum at a temperature of 4°C. After 4°C, it decreases, similar to any other normal behaviour liquid. Anomalous Expansion happens in the 0°C – 4°C range. In this, the density of water decreases, hence the volume of water increases. The density of water is maximum at 4°C with a value of 0.9998395 g/cm3 ~1g/cm3.
Reason for Anomalous Expansion of Water
Hydrogen bonding in water lead to some unusual and unique behaviour which is seen in daily life. Water shows highest density at 4oC and below this temperature its volume increase instead of decreasing. This anomalous behaviour water can be attributed by arrangement of hydrogen bonding among water molecules. The expansion of water is explained on the basis of ice-crystal in which molecules of water are further apart than in liquid water, thus ice occupies a large volume than the liquid water.
Ice has an “open” crystal structure which has a lot of space. Ice is less dense than water. When ice melts at 0°C, the water molecules lose this open structure and are more hydrogen bonding. This results in a lesser intermolecular distance between water molecules.
Hence, density increases from 0°C to 4°C. At 4°C, the density is maximum after 4°C, a rise in temperature results in increased intermolecular distance due to an increase in Kinetic Energy of the molecules.
hydrogen bonding in water above 4oC is temporary due to the high thermal energies. Thus water molecules can easily break their hydrogen bond and quickly reform new hydrogen bonds with some other molecules in the neighborhood.
When temperature falls below 4oC, water molecules start arranging themselves by lining up in such a manner that each water molecule can form up to maximum number of four hydrogen bonds. The increased number of hydrogen bonding results in the freezing of water into ice. In the form of ice water molecules are arranged in the more regular hexagonal patterns in the manner that empty spaces are created in the structure of ice and its volume expands up to 10%. The low density of ice can be seen when ice floats on water.
Consequences and Applications
Preservation of Aquatic Ecosystem
Anomalous expansion of water is important for sustaining aquatic life in cold regions or during winter. The liquids freeze from the bottom to upward but water freezes from the top to bottom.
During cold weather (winter), the top layer of a water body (lakes, rivers and the seas) cools first. The temperature of the top layer drops to 4°C. The top layer then becomes denser than other layers descends to the bottom of the water body. This process happens continuously to create a temperature gradient. The bottom most layer remains at 4°C. The surface water may freeze to form ice at its F.P. or even below the freezing point but the bottom layers remain practically at 4°C. At this 4°C, aquatic life can thrive. The temperature gradually decreases as the depth decreases. The topmost layer is the coldest. Eventually, the top layer freezes, thereby forming an insulating blanket that prevents further freezing of water underneath it to some extent. Thus, the entire body of water is rarely, if ever turned into ice. If there was no anomalous expansion, the waterworks have frozen altogether.
Weathering of Rocks
Water seeps into the cracks and crevices of rocks. During winter, when the temperature falls below 4°C, the water expands, resulting in hydrostatic pressure on the rock. This produces cracks on the rocks and helps in weathering the rock.
Pipeline Leakage and Bursting
Some water pipelines start to leak when the temperature falls below 4°C. This is due to the pressure created by the anomalous expansion of water.
In the cold regions, water supply pipeline burst if not properly insulated. The reason behind this is the anomalous expansion of water below 0oC.
As the water expands on freezing, bottle filled with water when placed in the freezer, will crack when water freezes. In cold weather, unless we put a chemical called glycol (antifreeze) in our car-radiator, the water will freeze and crack the radiator tubes for the same reason. Similarly, water supply pipes may also burst during severe winters due to the expansion of water on freezing.
Effect on State
hydrogen bonding affects on states of compounds. e.g. ethyl alcohol in which hydrogen bond is formed is a liquid while its isomer dimethyl ether lacking hydrogen bond is a gas at room temperature. Similarly unlike other hydrides of group VIA and VIIA, water and HF are liquids at 25°C in which strong hydrogen bonding is present.
Effect on Solubility
Hydrogen bond greatly increases the solubility of compounds. Various substances may mix and dissolve in each other if they have approximately the same type of polarity. e.g. ethyl alcohol is soluble in water as both have same type structure having polarized – OH group. Soap and detergent perform the cleaning action due to the polar part of their molecules are water-soluble due to H-bonding.
Effect on reactivity
Hydrogen bonding greatly increases the stability of compounds. Hydrogen bonding between molecules makes them less reactive. e.g. Strong hydrogen bonding in HF molecule makes it less reactive and weak acid which only 10% dissociates in aqueous solution as compared to other halogen acids which are 92% dissociated. [The order of reactivity of different halogen acids is HI > HBr > HCl > HF].
Low acidic strength of HF is due to HB. HF is a weak acid than HCl or HBr or HI due to HB in HF which entrapped H-atom tightly in between two electronegative atoms which decreases its ionization.
Hδ+–Fδ− ………… Hδ+–Fδ−
Effect on latent heat values
Hydrogen bonding greatly increases Latent heat of vaporization of liquids and latent heat of fusion of solids (as greater energy has to be spent in cleaving hydrogen bonding). e.g. water has high heat of vaporization of 40.6 kJ/mol and high heat of fusion of 6 kJ/mol due to presence of hydrogen bonding.
Effect on Bond Lengths and Bond Energy
Hydrogen bonding decreases bond length and increases bond energy of compounds.
The adhesive forces of paints and dyes is due to HB
Carbohydrates show HB due to O-H group in them
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