Covalent Bonding Complete SMTP

Chemical Bonding, Part IICovalent Bonding

Candidates should be able to

a. describe, including the use of 'dot-and-cross' diagrams, i. Covalent bonding, as in hydrogen; oxygen; chlorine; ii. Hydrogen chloride; carbon dioxide; methane and ethene *(ii) co-ordinate (dative

covalent) bonding, as in BF3.NH3

b. *explain the shapes of, and bond angles in, molecules by using the qualitative model of electron-pair repulsion (including lone pairs), using as simple examples BF3 (trigonal); CO2 (linear); CH4 (tetrahedral); NH3 (pyramidal); H2O( non-linear); SF6 (octahedral)

c. describe covalent bonding in terms of orbital overlap, giving and bonds

d. *explain the shape of, and bond angles in, the ethane, ethene and benzene molecules in terms of and" bonds

e. describe, interpret and/ or predict the effect of different types of bonding (ionic, covalent, hydrogen and intermolecular interactions) on the physical properties of substances

f. deduce the type of bonding present from given information

g. show an understanding of chemical reactions in terms of energy transfers associated with the making and breaking of chemical bonds.

* enhanced curriculum for SMTP

1. Molecules and Covalent Bonds

A covalent bond is the electrostatic force of attraction that two neighbouring nuclei have for a localised pair of electrons shared between them.

Covalent bonds are formed usually between non-metals and these elements are commonly from group IV, V, VI, VII. This arises because two non-metals both need to share electrons in order to achieve the stable noble gas configurations.

The bonding pair of electrons spends most of its time between the two atomic nuclei, thereby screening the positive charges from one another and enabling the nuclei to come closer together than if the bonding electrons were absent. The negative charge on the electron pair attracts both nuclei and holds them together in a covalent bond.

From an energy standpoint, when we say two atoms are chemically bonded we mean that the two atoms close together have lower energy and therefore are more stable than when separated. Energy is given off when atoms form a bond, and energy is taken in when bonds are broken.

When atoms combine by sharing electrons, molecules are formed. Examples of simple covalent molecules include:

Water (H2O) hydrogen molecule ( H2) Hydrogen chloride (HCl) methane (CH4)

1SYK/Sec 3 Chemistry (SMTP) 2012

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1.1 'Dot and Cross diagrams' of covalent molecules

The dot and cross structure of a covalent molecule shows how the valence electrons are arranged among the atoms in the molecule to show the connectivity of the atoms. The diagram shown does not indicate shapes of molecules.

Examples

a) Hydrogen chloride molecule (HCl):

Structural formula of hydrogen chloride:

b) Water molecule

Practice 1

Draw dot and cross diagrams to show the arrangement of electrons in each of the following molecules. Also include structural diagram.

Molecule Dot cross diagram Structural formula1. hydrogen

H-H

2. methane, CH4

3. ammonia

4. oxygen


Each chlorine and hydrogen atom contributes 1 electron for sharing. Hence the two atoms share two electrons so that each atom can achieve the stable noble gas configuration - the chlorine atom achieves the octet structure (8 valence electrons) while the hydrogen atom achieves duplet structure (2 electrons in the outer shell, just like helium)The two electrons that are shared constitute a single covalent bond.

Dot cross diagram: Structural formula:

xH H

CH

HH

H xxx

xC

H

H H

H

HH

H xxx

x

Nx N

HHH

O Ox xx

xxx O O

5. carbon dioxide

6. nitrogen

Exception to the octet rule

a) Odd electron molecules

i) Nitrogen oxide ii) Nitrogen dioxide

b) Molecules with incomplete octet (electron deficient)

i) BF3

c) Expanded octet ( Compounds with more than 8 electrons in the outershell per atom)

i) PCl5 ii) SF6

Note:1. Period 2 elements can only accommodate a maximum of 8 electrons in its outershell

and thus, cannot expand their octet.

2. Only elements of Period 3 and beyond can expand their octet to accommodate more than 8 electrons.


Cxx

xx OO C OO

N xxx Nx

x N N

After bonding with F, boron only has 6 electrons in its outer shell and hence is said to be electron deficient.

After bonding with Cl, P has 10 electrons in its outer shell.

After bonding with F, S has 12 electrons in its outer shell.

1.2 Co-ordinate (Dative Covalent ) Bonding

A covalent bond is formed by two atoms sharing a pair of electrons. The atoms are held together because the electron pair is attracted by both of the nuclei. In the formation of a simple covalent bond, each atom supplies one electron to the bond.

If both the electrons in a covalent bond come from only one of the atoms, the bond is called a co-ordinate bond or dative bond. Once the bond has formed, it is identical to any other covalent bond. It does not matter which atom the electrons come from.

In the formation of a dative bond: 1) one atom has a lone pair of electrons in the outer shell; Eg 1 nitrogen atom in ammonia has a lone pair of electrons.

Eg 2 Water molecule has 2 lone pairs of electrons.

2) The other atom has an empty orbital in the outer shell

3) The two orbitals overlap so that the electron pair is shared between the two atoms.

Example 1 - The ammonium ion, NH4+

The hydrogen ion in ammonium ion is electron deficient. Thus the empty 1s orbital of this hydrogen ion accepts a lone pair of electrons from the nitrogen atom, forming a dative/co-ordinate bond and acquiring a stable duplet configuration.

Example 2 - Aluminium chloride


Al2Cl6 dimer is formed because AlCl3 is electron deficient (two electrons short of a complete octet) and so, forms a dative bond with another AlCl3 by using its empty 3p orbital to accept a lone pair of electrons from Cl atom. In this way, aluminium atom would acquire a stable octet configuration.

Example 3 - BF3.NH3

In BF3, there are only six electrons around boron and boron is said to be electron deficient. Thus the boron atom can use its empty 2p orbital to accept a lone pair of electrons from nitrogen atom and so, forms a dative bond with NH3. In this way, the boron atom acquires the octet configuration and becomes stable.

1.3 Sigma and Pi bonds


Covalent bonds are formed when the orbitals of two neighbouring atoms overlap so that both nuclei attract the pairs of electrons between them. This can happen in two different ways making two different kinds of covalent bonds:

a) Sigma ( ) bond

A sigma bond is formed when the orbitals from two atoms overlap end-on (i.e. collinear or head-to head overlapping). Examples:

In a bond, the electron density is concentrated in the region between the two nuclei and is responsible for holding the atoms together against the mutual repulsion of the nuclei. It is a bond that lies along the internuclear axis.

All single bonds in molecules are sigma bond.

b) Pi () bonds

A pi bond is formed when the p orbitals of the two atoms overlap sideways (i.e. collateral or side-to- side overlapping). Electron density lies top and below the internuclear axis.

Diagram:

Before a pi bond is form, a sigma bond must be present. All molecules with double bonds will have one sigma bond and one pi bond. Molecules with triple bonds have one sigma and two pi bonds.


Internuclear axis

Internuclear axis

Internuclear axis

In a bond, there are two regions of electron density alongside the two nuclei (i.e. above and below the orbital , separated by a nodal plane - a region with zero electron density).

A pi bond is weaker than sigma bond because the overlapping of charge clouds is less than that in a sigma bond.

Examples: 1) ethane molecule, C2H6 2) ethene molecule, C2H4

3) benzene molecule, C6H6

In benzene, bond length of each bond is between a double bond and a single bond.Each bond is weaker than a double bond but stronger than a single bond.

1.4 Bond Energy, Bond Length and Bond Polarity

1.4.1 Bond energy

The strength of a covalent bond is measured by the bond energy. Bond energy is the energy required to break one mole of a covalent bond between two

atoms in the gaseous state. Examples:

H2 (g) H (g) + H (g) Bond energy = +436 kJ/mol

Br2 (g) Br (g) + Br (g) Bond energy = +190 kJ/molHCl (g) H (g) + Cl (g) Bond energy = +430 kJ/mol

The greater the bond energy, the stronger is the bond.

1.4.2 Bond Length


There are 7 sigma bonds in a molecule of ethane

The C-H bonds are sigma bonds.Between the two carbon atoms, there is one sigma bond and one pi bond

The six 2p orbitals of carbon overlap sideways and the electron cloud is above and below the plane of the ring, forming pi molecular orbitals all round the ring.The electrons within the pi orbitals can move throughout the ring and are thus delocalized.

The covalent bond length is the distance between the nuclei of two atoms joined by a covalent bond. Example: H-Cl bond length in hydrogen chloride

Cl-Cl Br-Br I-I C-C C=C CCBond length/nm 0.199 0.228 0.266 0.154 0.133 0.120Bond energy/kJ mol-1 244 193 151 350 610 840

From the table, the bond strength is in the order:

Cl-Cl > Br-Br > I-I and CC > C=C > C-C

Bond strength increases as bond multiplicity increases. This means that triple bond is stronger than a double bond, which is in turn stronger than a single bond provided that the two types of atoms in the bond are the same.

In general, the shorter the bond length, the stronger the covalent bond.

1.4.3 Bond Polarity

When a covalent bond is formed between two identical atoms, the electrons are equally shared and the bond is said to be non-polar.

Eg: H-H, Cl-Cl

Electron cloud is symmetrical due to equal sharing of electrons.

When a covalent bond is formed by two atoms of different electronegativities, a polar covalent bond is formed. Electronegativity is a measure of the attractive power of atoms for electrons in chemical bonds. It is a number between 0 and 4.0.

Na Mg Al Si P S Cl H O CElectronegativity 0.9 1.2 1.5 1.8 2.1 2.5 3.0 2.1 3.5 2.5

From the table, metals have low electronegativities and thus have small attraction for

electrons. Thus metals have a great tendency to lose electrons

Non-metals have high electronegativities and thus have a large attraction for electrons.

Thus non-metals have a large tendency to gain electrons.

Example : bond formed between hydrogen and chlorine


Bond length = r1 + r2

The more electronegative atom chlorine pulls the electrons of the covalent bond towards itself and so, acquire a small negative charge (-) while the other atom acquires a small positive charge (+). This unequal sharing of electrons is known as polarisation and the bond is known as polar covalent bond. Polar covalent bond has a dipole. (i.e there are two separate regions of charge).

Bond strength increases as differences in electronegativity increases.

Bond Difference in electronegativity Bond energy(kJ/mol)H-F 4.0 – 2.1 =1.9 565H-Cl 3.0 – 2.1 =0.9 430H-Br 2.8 – 2.1 = 0.7 360H-I 2.5 – 2.1 = 0.4 295

Difference in electronegativity between two atoms can be used to predict the type of bonding. A covalent bond is formed if the difference in electronegativity is small (usually < 1.5). An ionic bond is formed if the difference in electronegativity is large (usually > 1.5).

Practice 2

Use the electronegativity values on page 7, predict whether the following compounds would be ionic or covalent.a) sodium chloride 3.0-0.9 = 2.1 ionic

b) methane, CH4 2.5 – 2.1 = 0.4 covalent

c) aluminium chloride 3.0 – 1.5 covalent

d) magnesium oxide 3.5 – 1.2 = 2.3 ionic

e) sulfur dioxide 3.5 – 2.5 = 1.0 covalent

f) carbon dioxide 3.5 – 2.5 = 1.0 covalent

In general,

ionic compounds are usually formed between atoms of metals and non-metals. Exceptions are: AlCl3, SnCl4 and TiCl4.

Covalent compounds/molecules are formed between atoms of non-metals.

A alternative and better way to predict whether a compound is ionic or covalent is to study its melting and boiling point. Ionic compounds have high melting/boiling point while simple covalent molecules have low melting/boiling point.

Examples: Sublimation point of AlF3 = 1270°C. AlF3 is ionic.

Sublimation point of AlCl3 = 178°C. AlCl3 is covalent Tin (IV) chloride has melting point of -33°C and boiling point of 114°C. It is covalent.


A dipole moment µ is represented by an arrow with the tail at the positive centre and the head at the negative centre.

1.5 Shapes of Molecules

Shapes of molecules or ions that contain only non-metals can be predicted by the VSEPR Theory (Valence Shell Electron Pair Repulsion).

This theory states that:

1. The electron pairs distributed around a central atom will arrange themselves as far as possible in space to minimize their mutual repulsion. Electron pairs may be lone pairs (non-bonding pairs) or bond pairs.

Eg. NH3

2.Lone pair - Lone pair

electron repulsion > Lone pair - Bond pair electron repulsion > Bond pair – bond pair

electron repulsion

This is due to the lone pair of electrons being spread out more broadly than bond pairs and repulsions are greatest.

Different types of electron pair repulsions result in differences in bond angles.

Eg. Compare bond angles in CH4, NH3 and H2O

Determining Shapes of Molecules using VSEPR Theory

Step 1: Determine the central atom

Step 2: Draw the Lewis diagram (include the lone pairs)

Step 3: Count the number of electron pairs around the central atom; differentiate the bond pairs and lone pairs.

Step 4: Classify the general shape of the molecule by the total number of electron


pairs

Step 5: Determine the actual shape of molecule (refer to table on pg 10)

Total no. of electron

pairs

General Shape(Step 4)

No. of bond pairs of

electrons

No. of lone pairs of

electrons

Shape of molecule (Step 5)

2 Linear 2 0 Linear3 Trigonal Planar 2 1 Bent

3 0 Trigonal Planar4 Tetrahedral 2 2 Bent

3 1 Trigonal Pyramidal4 0 Tetrahedral

5 Trigonal Bipyramidal

2 3 Linear3 2 T-shape4 1 See-saw5 0 Trigonal Bipyramidal

6 Octahedral 2 4 Linear3 3 T-shape4 2 Square planar5 1 Square pyramidal6 0 Octahedral

Practice 3

Using VSEPR Theory, determine the shapes of the following:

Molecule Sketch of Lewis Structure

Total no. ofelectron

pairs around central atom

No. ofbond pairs

around central atom

No. oflone pairs

around central atom

Shape

BF3 3 3 0 Trigonal planar

CO2 2 2 0 linear

CH4 4 4 0 tetrahedral

NH3 4 3 1 Trigonal pyramidal


H2O 4 2 2 Bent

SF6 6 0 6 octahedral

1.6 Molecule Polarity

A molecule is polar and thus, has a dipole moment () if 1. its bonds are polarised, and

2. it is not symmetrical (i.e the individual dipoles do not cancel each other out).

Examples:

1) H-F 2) CO2 3) H2O

Polar The two dipoles cancel out; non-polar

4) CCl4 5) CHCl3

1.7 Structure and Properties of Simple Covalent Molecules

Simple covalent molecules are said to have simple molecular structures. Examples: iodine


Structure of Iodine solid

polar

Structure of molecular substances Intermolecular forces are defined to be the electrostatic interactions between atoms and

molecules that act to hold the particles together. The forces between the molecules , intermolecular forces, are weak but the covalent bonds between the atoms are strong.

Example: covalent bonds:150-1100 kJ/mol

Intermolecular forces: 1-50 kJ/mol

When covalent molecules boil, only the weak intermolecular forces are overcome, not the covalent bonds.

Physical Properties of simple covalent molecules:

i. have low melting and boiling points (boiling points of simple covalent compounds are usually below 200C). Simple covalent substances (E.g H2O, NH3, O2, CO2 ) exists as simple discrete covalent molecules. Only a small amount of energy is required to overcome the weak intermolecular forces , hence the boiling points of simple covalent molecules are low.

Liquids with low boiling point and vaporise easily at room temperature are said to be volatile.

ii. do not conduct electricity in any states. Why?

Simple covalent substances exist as molecules. They are not charged and do not have

mobile ions. Also there are no mobile electrons as all the electrons are either used up

to form covalent bonds or are located in electron shells.

iii. usually insoluble in water but soluble in organic solvents.

Exceptions:Gases such as HCl, SO2, NH3 can dissolve in water because of chemical reactions.

E.g.

Sugar and ethanol are soluble in water because of their ability to form hydrogen bonds with water molecules.

2 INTERMOLECULAR FORCES

Electrostatic interactions between atoms and molecules that act to bind the particles together are intermolecular forces.

Intermolecular forces are weaker than chemical bonds. Chemical bonds: 150-1100 kJ/mol Intermolecular forces: 1-50 kJ/mol

2.1 Types of Intermolecular Forces

a) Hydrogen bond


Electrostatic attraction between a hydrogen bonded to nitrogen, oxygen or fluorine and the lone pair of a neighbouring nitrogen, oxygen or fluorine atom is called a hydrogen bond.

Bond Strength of hydrogen bond : H-F > H2O > NH3 as F is most electronegative, followed by O, then N

Boiling points : NH3 ; -33.3oC < HF ; 19.7oC < H2O ; 100oC

Notice that in each of these molecules;

i. The hydrogen is attached directly to one of the most electronegative elements (N, O, F) causing the hydrogen to acquire a significant amount of positive charge.

ii. Each of the elements (X = N, O, F) to which the hydrogen is attached also has at least one “active” lone pair.H-X has a very polar bond due to the large electronegativity difference. When X is bonded to H they pull the bonding pair well away from H, making H positive and poorly shielded. H carries the δ+ and X the δ-.

iii. Lone pairs at level n = 2 (2p orbitals) have the electrons contained in a relatively small volume of space which therefore has a high density of negative charge.

Boiling points of hydrogen bonded molecules

HF, H2O and NH3 have higher than expected melting points and boiling points compared with the hydrides


NH3 forms an average of 1 hydrogen bond per molecule

H2O forms an average of 2 hydrogen bonds per molecule

HF forms an average of 1 hydrogen bond per molecule

of other elements in their respective Group.

b) Van der Waals' forces

Arises from induced dipole or permanent dipole.

i) Permanent dipole-permanent dipole interaction

Permanent dipole-permanent dipole forces exist between two molecules with permanent dipole moments. Negative end of one polar molecule attracts the positive end of another

e.g. CH3Cl have net dipole moment.

ii) Induced dipole or London dispersion forces

The second type of van der Waals' forces is the instantaneous dipole-induced dipole forces or dispersion forces. It is a weak, short range electrostatic force of attraction between non-polar molecules with temporary dipoles . It arises from the weak attractive force of the electrons on one molecule for the nuclei of another molecule.

Cl2(s) Cl2(g) 25 kJ/mol (dispersion force broken)Cl2(g) 2Cl(g) 244 kJ/mol (covalent bond broken)

Origin of London dispersion forces (VDW)

electrons are in constant motion instantaneous dipole moment forms when there are more electrons on one

side of the molecule


Induced temporary dipole

Original temporary dipole

- +An instantaneous dipole

- + - +

instantaneous dipole moment induces an instantaneous dipole moment in a neighbouring molecule. The result is that particles will then attract each other

Larger atom has a greater electrostatic attraction (between electrons of one atom and the nuclei of an adjacent atom)

Weak dispersion forces are always present in all molecules.

For non-polar molecules, these are the only forces present.

e.g. argon, only instantaneous dipole-induced dipole or dispersion forces

e.g. HCl molecules has permanent dipole-permanent dipole and instantaneous dipole-induced dipole interactions.

Factors influencing the strength of dispersion forces:

i) Size (number of electrons)

Boiling point is used to indicate the strength of intermolecular forces.In boiling only van der Waals’ forces are broken (for substances made of simple, discrete molecules only)

Molecule No of electrons b.p. / oCH2 2 - 253Ne 10 - 246N2 14 - 196Ar 18 - 186Cl2 34 - 35

Strength of van der Waals’ forces increases with the number of electrons contained in a molecule.

The greater the electron cloud, the greater the magnitude of momentary distortion (polarisability of an atom or molecule) and the greater the instantaneous dipole-induced dipole interaction/dispersion forces.

ii) Shape of molecules

Molecule b.p. / oC densityn-pentane (C5H12) 36 0.626

2-methylbutane (C5H12) 28 0.620


Array of molecules which havetemporary dipole

2,2-dimethylpropane (C5H12) 9 0.591

For molecules with similar molecular formula but different structural formula, the boiling point of straight chain molecules are higher than branched molecules.

Attraction between linear molecules is stronger than between spherical molecules because of larger surface area of contact between the molecules and thus greater dispersion of electrons and greater the strength of induced dipole-induced dipole interaction.

Note:

Instantaneous dipole-induced dipole interaction or dispersion forces are present in all systems, but only become more significant if there are no stronger intermolecular forces present to mask their effect.

2.2 Comparing the strength of different types of intermolecular forces

In general,

Hydrogen bonds is usually stronger than van der waals forces (permanent dipole – permanent dipole or instantaneous dipole-induced dipole)

3. Giant Covalent/Molecular Structures (Macromolecules)

A macromolecule is a giant molecule made by large number of atoms covalently bonded together.Examples of macromolecules include diamond, graphite, silicon, silica, and poly(ethene).

3.1 Diamond

Diamond and graphite are allotropes of carbon. Allotropes are different forms of the same element.


Structure of diamondIn diamond, each carbon atom is joined to four others by strong covalent bond. This is called the tetrahedral arrangement of atoms. The tetrahedral network structure is repeated and forms a giant molecule known as macromolecule.

Deducing properties of diamond from its structure

1. Diamond is the hardest substance known because a large amount of energy is needed to break the strong covalent bonds operating in 3-dimensions.

2. Diamond has a very high melting point (3500°C) and boiling point (4800°C) . This is because diamond contains millions of carbon atoms joined by strong covalent bonds. A large amount of energy is needed to break the strong covalent bonds throughout the structure before melting occurs.

3. Diamond does not conduct electricity. In diamond, all the four valence electrons are used to form covalent bonds. Hence there are no mobile electrons that move through the structure.

4. Diamond is not soluble in water because the attractions between solvent molecules and carbon atoms are not strong enough to overcome the strong covalent bonds in diamond.

5. Diamond is transparent.

Uses of diamond

Diamonds are used for gemstones in jewellery

Because diamond is very hard, they are used as tips for drills and other cutting tools. Diamonds are used for drilling, grinding and polishing very hard surfaces.

3.2 Graphite

Deducing properties of graphite from structure

1. Graphite has a very high melting point and boiling point. The covalent bonds between the carbon atoms within each layer are strong and require a lot of energy to break it.

2. Graphite is soft and slippery. The layers of atoms are held by weak van der Waals' forces, allowing the layers of atoms to slide over one another, hence causing graphite to be soft and slippery.

3. A carbon atom has four outer shell electrons. In graphite, each carbon atom only uses three valence electrons to form covalent bond, leaving one electron that is free. These free electrons are said to be delocalised as they can move along the layers from one carbon atom to the next when graphite is connected to a battery. This explains why graphite can conduct electricity.

4. Graphite is black.

5. Graphite is insoluble in water, for the same reason as diamond.


Graphite is made of flat layers of carbon atoms.

Within each layer, each carbon atom is bonded to three carbon atoms by strong covalent bonds. This forms rings of six carbon atoms that join together to form two-dimensional flat layers.

The layers of carbon atoms lie one on top of the other and are held by weak forces known as van der Waals forces.

Uses of graphite

Used as a dry lubricant as it is slippery, thermally stable unlike oil and does not attack rubber.

Graphite is baked with clay and made into pencil lead. As it is soft, the carbon layers flake off and stick to the paper when we write.

Because of graphite is a good conductor of electricity, it is used to make inert electrodes for batteries and for electrolysis. It is also used as brushes for electric motors.

3.3 Silicon dioxide or silicon (IV) oxide, SiO2

3.4 Silicon

3.5 Poly(ethene) (self-study section)

4. Comparing covalent substances with different structures (Self-study)

Type of structure Typical physical propertiesSimple molecular structure

Low melting and boiling points, that is, volatile; usually exist as liquids or gases at room temperature.


Silicon (IV)oxide is commonly called silica. It is the main component in sand.

Each silicon atom is bonded to four oxygen atoms in a tetrahedral structure and each oxygen atom is bonded to two silicon atoms.

Silicon (IV) oxide is a hard solid, has high melting and boiling points and does not conduct electricity.

Powdered silicon dioxide is a mild abrasive used in toothpaste.

Silicon has a giant molecular structure similar to diamond. Each silicon atom is bonded to four other silicon atoms in a

tetrahedral arrangement. Properties of silicon are similar to those of diamond. Silicon

is a hard substance and has very high melting and boiling points. Pure silicon does not conduct electricity but becomes a semiconductor when impregnated with impurities.

Poly(ethene), a polymer is the simplest form of plastic.

It contains thousands of carbon and hydrogen atoms joined together to form a long chain. Each carbon atom is surrounded tetrahedrally by two hydrogen atoms and two other carbon atoms.

Poly(ethene) has a high melting point, can be easily moulded by heat, and does not conduct electricity.

(e.g. hydrogen, water, ammonia, iodine)

Not able to conduct electricity in all states. Insoluble in water but usually soluble in organic solvents.

Giant molecular /covalent structure(e.g. diamond, graphite, silica)

High melting and boiling points, that is, non-volatile, exist as solids at room temperature.

Not able to conduct electricity except graphite Insoluble in all solvents.

5. Differences between ionic and covalent compounds (Self-study):

Ionic Compound Covalent Compound (with simple molecular structure)

Melting/Boiling points High (non-volatile) Low (volatile)Electrical conductivity conduct electricity in aqueous or

molten state. Do not conduct when solid.

Do not conduct electricity at all.

Solubility in water Usually soluble Usually insolubleSolubility in organic solvent

Usually insoluble Usually soluble

6. Influence and importance of H-bonding (Self-study)

The structure and property of water and iceIn the ice lattice, the O-H hydrogen bond (0.180 nm) is much longer than the covalent O-H bond (0.096 nm). This results in ice having an open but rigid 3D tetrahedral structure which holds the water molecules further apart than in the liquid form (hence the larger volume and lower density of ice)

http://www.tms.org/pubs/journals/JOM/9902/Schulson-9902.html

Solubility of polar substances in waterThe ability of other molecules capable of hydrogen bonding will increase the solubility of that compound in water. For example, ethanol and sugar are very soluble in water as their molecules are able to form hydrogen bonds with water molecules. Also, ionic compounds can also dissolve in water due to the favourable ion-dipole interaction.

However, non-polar substances will not dissolve in water. It will result in the formation of an immiscible layer. Solubility of one substance in another is dependent on the forces of attraction between units of the substance. Substances with similar intermolecular attraction would dissolve in one another.

High boiling point of water (100°C) due to the presence of hydrogen bondingThis is of particular importance because without the operation of hydrogen bonds, water would be a gas under normal atmospheric conditions; oceans, lakes and rivers would never exist and it would never rain.

Dimerization of carboxylic acids


http://www.tms.org/pubs/journals/JOM/9902/Schulson-9902.html

Protein folding in DNA


Covalent Bonding Complete SMTP

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