Chapter 3Chemical Bonding

Metals generally have a tendency to lose electrons to form positive ions called cations; this is their electropositive character. Usually metals have low values of ionisation energy, so their atoms lose their outer electrons easily. Alkali metals (Na, K) are highly electropositive; alkaline earth metals (Mg, Ca) lose their outer electrons less easily.

Non-metals have an affinity towards electrons. They tend to gain electrons and become negatively charged ions called anions; they are electronegative elements. Fluorine is the most electronegative element, followed by oxygen.

Covalent compounds mostly exist as discrete neutral molecules. The forces between these molecules (intermolecular forces) are very weak compared with the bonding forces inside them. Lower molecular mass covalent compounds are gases or low boiling liquids because little energy is needed to separate their molecules. Examples: CH₄, NH₃, H₂S, HCl, CO₂ and SO₂ are gases at room temperature.

Diamond (carbon). Carbon, phosphorus and sulphur exist as covalent solids, in both amorphous and crystalline forms. In diamond, each carbon atom is surrounded by four other carbon atoms linked together by strong covalent bonds. Graphite and rhombic sulphur are other examples.

In metals the atoms are strongly held and arranged in the form of rows one above the other. When pressure is applied on the metals, the upper rows of atoms slip past the lower rows. As a result, their shapes are changed. Metals can therefore be easily drawn into wires and hammered into sheets. The mobile electrons keep holding the rows together, so the metal bends instead of breaking.

Yes. A coordinate covalent bond is a type of covalent bond, and once formed it is as strong as an ordinary covalent bond. The difference between a covalent bond and a coordinate covalent bond lies in the way they are formed. Once such bonds are formed, there does not remain any difference. All four N–H bonds in the ammonium ion behave exactly alike.

H shares one pair with O; that O shares one pair with N; N shares two pairs with a second O (double bond) and gives its lone pair to the third O (coordinate bond). Every O and N then has an octet, and H has a duplet.

Ionic bonds are formed between elements that lose electrons easily (metals) and elements that gain electrons easily (non-metals). Alkali and alkaline earth metals are electropositive metals which form bonds with electronegative elements like oxygen and chlorine. Examples: NaCl, KCl, MgF₂, CaCl₂.

Non-metals, which combine with other non-metals to form a wide variety of molecular substances: H, C, N, O, S, P and the halogens. They lower their energy by mutually sharing their electrons. Examples: H₂, Cl₂, H₂O, CH₄, NH₃.

A coordinate bond is formed when a species has an electron pair to donate to another species.

(a) Neither BF₃ nor AlCl₃ has a lone pair on its central atom; B and Al are both acceptors, so no ordinary coordinate bond forms between them. AlCl₃ forms one with a donor such as NH₃ (H₃N→AlCl₃), or with a Cl lone pair of another AlCl₃ molecule (Al₂Cl₆).

(b) A lone pair on the oxygen of dimethyl ether bonds the proton: (CH₃)₂O + H⁺ → [(CH₃)₂OH]⁺.

A coordinate bond needs a species that has an electron pair to donate and one that can accept it. A compound with no lone pair and no empty space cannot form one, such as methane, CH₄: carbon has four bond pairs and no lone pair, and hydrogen already has its duplet.

Metal atoms, which usually have low values of ionisation energy, so they lose their outer electrons easily (Na, Mg, Al, Cu, Fe). The nuclei of metallic atoms cannot hold their outer electrons firmly, and these electrons become a sea of mobile electrons.

Atoms have a tendency to decrease their energy. They can do this by combining with other atoms. It is a natural phenomenon because it increases the stability of atoms. Atoms form bonds because they would like to lower their energy by completing their duplet or octet.

Helium has two electrons in its outermost shell, while all other noble gases have eight, and these gases are unreactive. Having two electrons (for hydrogen and helium, which have only the first shell) or eight electrons in the outermost shell means stability and hence unreactivity. This principle was named the duplet or octet rule.

Neither their atoms combine with themselves nor with other atoms. The probable reason for this lack of reactivity was their stability, due to the presence of two electrons in helium and eight electrons in the outermost shells of the rest of the gases.

For a sodium atom it is easier to lose one electron and stabilise itself than to gain seven electrons to complete its octet. Sodium therefore adopts the energetically easier path and loses its electron to form a bond.

It is energetically favourable for a hydrogen atom to lose one electron to become a proton (H⁺) or gain one electron to become a hydride ion (H⁻). In the latter case, it completes its duplet. It can also complete its duplet by sharing (H₂, HCl).

The arrangement of electrons around the nucleus of an atom in shells and sub-shells is called electronic configuration. Example: Na 2, 8, 1; Cl 2, 8, 7.

A chemical bond is a force of attraction between atoms which holds them together in the form of a molecule or a compound.

When atoms approach each other there are two possibilities: they may attract or repel each other. If the forces of attraction dominate the forces of repulsion, the energy of the system gets lowered and the two atoms react to form a new molecule. Conversely, the two atoms simply move away from each other.

(1) Ionic bond, (2) covalent bond, (3) coordinate covalent bond, and the metallic bond found in metals.

An ionic bond is a bond which is formed by the complete transference of electron or electrons from one atom to another atom. The oppositely charged ions formed are held together by the electrostatic force of attraction. It is also called an electrovalent bond, and compounds having such a bond are called ionic compounds.

An electron from the outermost shell of the sodium atom is transferred to the outermost shell of the chlorine atom, and in doing so both atoms acquire the electronic configurations of their nearest noble gases. The Na⁺ and Cl⁻ ions are held together by electrostatic attraction.

Calcium, an alkaline earth metal, loses two electrons, one to each of two chlorine atoms, to form calcium chloride:

The electron or electrons which take part in a chemical reaction come only from the outermost shells of the atoms.

The ions of an ionic compound surround each other three-dimensionally to form a crystal lattice. Since ions are spherical and oppositely charged, they surround each other from all sides. In NaCl each Na⁺ is surrounded by six Cl⁻ ions.

Since ions are spherical and oppositely charged, they can surround each other from all the sides; therefore ionic bonds are non-directional.

KCl, MgF₂, NaF, KBr, CaF₂, NaCl and CaCl₂.

A covalent bond is a bond formed by the mutual sharing of an electron pair provided by the bonded atoms. This is called a single covalent bond. It is the strong electrostatic attraction between the shared electrons and the two nuclei.

In some compounds the atoms share two electrons each to form a double covalent bond, and in the same way atoms can share three electrons each to form a triple covalent bond. Double and triple bonds have two and three electron pairs mutually shared between the two atoms. A single bond is shown by a single line (−), a double bond by two lines (=) and a triple bond by three lines (≡).

When two atoms approach, the electrons of one atom come under the attractive influence of the nucleus of the other atom; this lowers the energy. The electrons and nucleus of one atom also repel those of the other, which increases the energy. The two atoms bring themselves to such a distance that the attractive forces dominate the repulsive forces. The total energy at this distance will be minimum and thus a stable molecule is formed.

The force of attraction of an atom for the shared pair of electrons is called its electronegativity. Fluorine is the most electronegative element in the periodic table, followed by oxygen.

When a covalent bond is formed between two identical atoms, the shared pair of electrons lies at the middle of the two bonded atoms. This gives rise to a pure covalent bond which is essentially non-polar (H₂, Cl₂).

When a covalent bond is formed between two non-identical atoms, the shared pair of electrons bends towards the atom which has more electronegativity. A partial negative charge is created on the more electronegative atom and an equal partial positive charge on the less electronegative atom: a polar covalent bond (HCl, H₂O).

In the HCl molecule the shared pair of electrons bends towards the more electronegative chlorine atom, creating a partial negative charge on it. H gets an equal partial positive charge, so the H–Cl bond is a polar covalent bond: H(δ+)–Cl(δ−).

H₂O has 2 lone pairs on oxygen, NH₃ has 1 lone pair on nitrogen, and CH₄ has none.

Coordinate covalent bond is a type of covalent bond in which the shared electron pair is donated by one atom only. This bond is formed when a species has an electron pair to donate to another species. The species which donates the electron pair is called a donor, while that which accepts it is called an acceptor. An arrowhead (→) pointing towards the acceptor represents this bond.

Acids provide protons (H⁺) when dissolved in water. This proton has an empty outer shell and can accept one of the two pairs of electrons present on the oxygen atom of the water molecule. As a result, a hydroxonium ion (H₃O⁺) is formed. The positive charge covers the whole ion, and all three bonds of oxygen behave exactly alike.

During the reaction, an electron pair from nitrogen of ammonia fills the partially empty outer shell of boron present in boron trifluoride, forming a coordinate covalent bond: H₃N→BF₃.

When a coordinate covalent bond breaks, both the bonded electrons go to the donor species.

In ammonium chloride a coordinate covalent bond links the nitrogen of ammonia and the proton. The positive charge is spread all over the ammonium ion, and all the four bonds between nitrogen and hydrogen behave exactly alike. This proves that the difference lies only in the way they are formed; once formed, there does not remain any difference.

A metallic bond is a type of chemical bond which has positively charged ions bound together by the mobile electrons. A metal appears to have a sea of electrons in which all the nuclei of the atoms are submerged.

The strength of a metallic bond depends upon two factors: the number of positive charges present on the positive ions and the number of mobile electrons set free by each atom.

In sodium metal each sodium atom sets free only one electron, so the metallic bond is not very strong. In magnesium, each atom releases two electrons to acquire two positive charges, so the metallic bond is stronger. This explains why magnesium melts at a higher temperature than sodium (650 °C against 98 °C).

The presence of freely moving electrons in metals makes them good conductors of heat and electricity.

In metals, the atoms are strongly held and arranged in the form of rows one above the other. This arrangement makes them hard and heavy.

Metals are extensively used in machinery, automobiles, railways, aircraft, rockets, the construction industry, the electronics industry, jewellery, electric wires and many more.

Metals generally have a tendency to lose electrons to form positive ions called cations. This property is called the electropositive character of metals. It is related to reactivity: metals which lose electrons easily are more reactive.

Alkali metals (Na, K) are highly electropositive elements and undergo reactions very easily; sodium and potassium react vigorously with water and halogens to give hydroxides and halides, and with acids to give salts and hydrogen. Alkaline earth metals (Mg, Ca) lose their outer electrons less easily, so they are less electropositive and their reactions are less vigorous.

Non-metals have an affinity towards electrons. They tend to gain electrons and become negatively charged ions called anions; they are therefore named electronegative elements. Non-metals react with metals to form ionic bonds, and with other non-metals to form a wide variety of molecular substances.

Fluorine is the most electronegative element in the periodic table, followed by oxygen, then chlorine and nitrogen.

The forces of attraction which are present between the molecules of elements and compounds are named intermolecular forces of attraction. They are generally very weak compared with the bonding forces between atoms; among the three states they are weakest in gases and strongest in solids.

The melting and boiling points of substances depend on the strength of these forces. The stronger the forces among the molecules of a liquid, the higher is its boiling point and vice versa; similarly, the stronger the forces, the higher the melting point of a solid.

These attractive forces are present between the molecules of a polar compound like HCl. Chlorine attracts the shared pair more strongly, so the bond is polar (H(δ+)–Cl(δ−)). Due to these partial charges the molecules of HCl start attracting each other. These forces of attraction are called dipole-dipole forces. Such compounds show relatively higher melting and boiling points.

Hydrogen bonding is a special case of dipole-dipole attractive forces. When hydrogen is covalently bonded to highly electronegative elements like F, O or N, the large difference of electronegativity makes the bond highly polar, and strong dipole-dipole attractions are observed among the molecules. The attractive force between the molecules of water is called hydrogen bonding.

The strength of the hydrogen bonds causes water to have relatively higher melting and boiling points as compared to compounds like H₂S and NH₃. Water boils at 100 °C, NH₃ at −33 °C and H₂S at −60 °C.

In ionic compounds the oppositely charged ions are held together by the strong electrostatic force of attraction in the form of a crystal lattice. Since the ions are rigid, such compounds exist as very stable solids with significantly high melting points. The melting point of NaCl is 801 °C.

If an external force is applied on the crystal lattice, it breaks easily. It shows that ionic solids are highly brittle. The force pushes like charges next to each other and they repel.

In the solid form, ionic compounds do not conduct electricity because the ions are tightly held and cannot move. In the molten state the ions get free and start conducting electricity. In aqueous solution the free ions also move towards their electrodes.

In the process of hydration, water molecules surround and interact with ions or molecules. Water not only breaks the electrostatic attraction in an ionic solid but also hydrates the free ions.

When we mix two solutions of ionic compounds, the positive ions of one compound may react with the negative ions of the other to form a new compound. A white precipitate of silver chloride comes out of the aqueous solution.

Electrolysis of molten sodium chloride gives sodium metal and chlorine gas. Electrolysis of aqueous sodium chloride gives sodium hydroxide and chlorine gas.

Elements on the right side of the periodic table (except inert gases) exist as covalently bonded diatomic molecules. Due to very weak forces of attraction between their molecules, their densities and boiling points are very low. Bromine (Br₂) exists as a volatile fuming liquid.

In diamond, each carbon atom is surrounded by four other carbon atoms linked together by strong covalent bonds. Due to this rigid structure, diamond is the hardest thing on this planet. It is used as a cutting, polishing and drilling tool.

Graphite consists of a layered structure made of hexagonal rings of carbon. Since the layers are not bonded strongly, they can slip past each other, so graphite is used as a lubricant in industry. The layers have mobile electrons between them, so graphite is a good conductor of electricity and is used as an electrode.

Coal is the amorphous form of carbon, whereas diamond and graphite are its crystalline forms. Coal is used as a fuel in electricity generating plants.

Binary covalent compounds generally exist as low boiling gases, except water. Methane (CH₄), ammonia (NH₃), hydrogen sulphide (H₂S), hydrogen chloride (HCl), nitrogen dioxide (NO₂), carbon dioxide (CO₂) and sulphur dioxide are all gases at room temperature, because the forces between their molecules are weak.

Covalent molecules like hydrogen chloride, sulphuric acid and nitric acid ionise completely in water, behaving as very strong acids.

When hydrogen is covalently bonded to highly electronegative elements like F, O or N, the bond becomes highly polar and strong dipole-dipole attractions (hydrogen bonds) are observed among the molecules. Fluorine is the most electronegative element and very small, so HF molecules are held by strong hydrogen bonds. HCl has only weaker dipole-dipole forces.

So HF boils at 19.5 °C while HCl boils at −85 °C. The book notes that water and hydrogen fluoride are liquids at room temperature.

Covalent compounds are usually insoluble in water but soluble in non-polar solvents like ether and benzene. Water molecules are polar and held to each other by hydrogen bonds. Most covalent compounds are non-polar: their molecules cannot attract water molecules strongly enough to break those hydrogen bonds and mix in ("like dissolves like").

Exceptions are polar covalent compounds that hydrogen-bond with water or react with it, such as sugar, ethanol, ammonia and HCl.

The presence of freely moving electrons in metals makes them good conductors of heat and electricity.

The mobile electrons at the hot end gain kinetic energy and move very fast through the metal, colliding with ions and other electrons and passing the energy on. The vibrating ions also pass energy to their neighbours. Because the free electrons travel so quickly, metals conduct heat much better than non-metals.

Non-metals combine with other non-metals to form a wide variety of molecular substances. Nitrogen forms six common oxides with oxygen:

All are covalent; nitrogen bonds to oxygen by sharing electrons.

Ionic compounds generally react in an aqueous solution. When we mix two solutions of ionic compounds, the positive ions of one compound may react with the negative ions of the other to form a new compound.

Ag⁺ and Br⁻ combine to form silver bromide, which is insoluble and comes out as a pale yellow (cream) precipitate. Sodium nitrate stays dissolved.

Lower molecular mass covalent compounds are gases or low boiling liquids; high molecular mass covalent substances exist as solids. Cl₂ and I₂ are both non-polar diatomic molecules held by weak intermolecular forces, but these forces (London or dispersion forces) grow with the size of the molecule.

An I₂ molecule is much bigger, with 106 electrons against 34 in Cl₂, so the forces between I₂ molecules are much stronger. Iodine therefore melts only at 114 °C, while chlorine boils at −34 °C.

Ionic bond

This chemical bond is formed as a result of the tendency of atoms to lose or gain electron or electrons to acquire the electronic configuration of the nearest noble gas, because this is a more stable electronic structure.

Formation of sodium chloride

Sodium chloride is formed when the elements sodium and chlorine react chemically. Their electronic configurations are Na 2, 8, 1 and Cl 2, 8, 7. An electron from the outermost shell of the sodium atom is transferred to the outermost shell of the chlorine atom, and in doing so both these atoms acquire the electronic configurations of their nearest noble gases.

The electrons which take part in a chemical reaction come only from the outermost shells. Sodium chloride contains the positively charged sodium ions (Na⁺) and negatively charged chloride ions (Cl⁻). These oppositely charged ions are held together by the electrostatic force of attraction. The chemical bond thus formed is called an ionic or electrovalent bond. Calcium, similarly, loses two electrons to form CaCl₂. An ionic bond is therefore a bond formed by the complete transference of electron or electrons from one atom to another atom.

Covalent bond

Some atoms decrease their energy by mutually sharing their electrons. When two atoms approach each other, the electrons of one atom come under the attractive influence of the nucleus of the other atom; this new force of attraction lowers the energy. The electrons and the nucleus of one atom also repel those of the other atom, which increases the energy. The two atoms bring themselves to such a distance that the attractive forces dominate the repulsive forces. The total energy at this distance is minimum and thus a stable molecule is formed.

A covalent bond is therefore a bond formed by the mutual sharing of an electron pair provided by the bonded atoms. This is called a single covalent bond. Atoms may share two electrons each to form a double covalent bond, or three each to form a triple covalent bond.

Sodium chloride contains Na⁺ and Cl⁻ ions held together by the electrostatic force of attraction. These ions then surround each other three-dimensionally to form a crystal lattice.

The NaCl lattice

Properties that follow

In ionic compounds the ions are held together by strong electrostatic forces in the form of a crystal lattice. Since the ions are rigid, such compounds exist as very stable solids with significantly high melting points (NaCl melts at 801 °C). If an external force is applied on the crystal lattice, it breaks easily: ionic solids are highly brittle.

Nature of the metallic bond

Usually metals have low values of ionisation energy. Their atoms can therefore lose their outer electrons easily. In sodium metal each sodium atom is surrounded by eight other sodium atoms. The outer electrons of these atoms move freely between the vacant spaces. No electron remains attached to any particular nucleus; instead all the electrons at the same time get attached to all the nuclei. When all the atoms attract all the electrons collectively, they are bound together. A metal appears to have a sea of electrons in which all the nuclei are submerged.

1. High melting and boiling points

The strength of a metallic bond depends on the number of positive charges on the ions and the number of mobile electrons set free by each atom. Each sodium atom sets free only one electron, so its metallic bond is not very strong; each magnesium atom releases two electrons, so magnesium melts at a higher temperature (650 °C against 98 °C).

2. Conduction of heat and electricity

The presence of freely moving electrons in metals makes them good conductors of heat and electricity.

3. Hard and heavy

In metals the atoms are strongly held and arranged in the form of rows one above the other. This arrangement makes them hard and heavy.

4. Malleable and ductile

When pressure is applied on metals, the upper rows of atoms slip past the lower rows. As a result their shapes are changed. Metals can therefore be easily drawn into wires and sheets.

5. Lustre

Metals usually show metallic lustre: the free electrons absorb and re-emit the light falling on the surface, so metals shine.

Properties of ionic and covalent compounds are adequately explained on the basis of the type of attractive forces present between them.

Structure of graphite

Graphite consists of a layered structure, made of hexagonal rings of carbon. In a layer each carbon atom forms covalent bonds with only three other carbon atoms, using three of its four outer electrons.

Mobile electrons

These layers in graphite have mobile electrons in between them. The fourth outer electron of each carbon atom is not fixed in a bond; these electrons are delocalised and free to move along the layers, like the electrons in a metal.

Conduction

When a voltage is applied, the mobile electrons drift through the layers and carry the current. Graphite is therefore a good conductor of electricity and is also used as an electrode, in battery cells and in electric arc furnaces to produce steel.

In metals, the atoms are strongly held and arranged in the form of rows one above the other. This arrangement makes them hard and heavy.

Hard

Every positive ion is attracted by the whole sea of mobile electrons. In sodium each atom is surrounded by eight others; these strong, all-round attractions hold the rows firmly, so most metals are hard.

Heavy

Because the atoms are packed so closely, many fairly heavy atoms fit into a small volume: metals have high densities, such as iron 7.9 g cm⁻³ and copper 8.9 g cm⁻³.

Exceptions

Sodium and potassium set free only one electron each, so their metallic bonding is weak: they are soft enough to cut with a knife and light enough to float on water (densities 0.97 and 0.86 g cm⁻³).

Why AlCl₃ joins up

A coordinate covalent bond is formed when a species has an electron pair to donate to another species. In AlCl₃ aluminium shares its three outer electrons with three chlorine atoms, so it has only six electrons in its outer shell: like boron in BF₃, it has a partially empty outer shell and can accept an electron pair. Each Cl atom has three lone pairs to donate.

The bridge

A giant covalent structure

Compounds may exist as a network arrangement of their atoms, for example covalent compounds like sand (SiO₂). Sand is mostly silicon dioxide (silica, as quartz). It is not made of small SiO₂ molecules: it is one giant network of atoms joined by covalent bonds, like diamond.

Properties that follow

Why atoms form bonds

Atoms have a tendency to decrease their energy. They can do this by combining with other atoms. It is a natural phenomenon because it increases the stability of atoms.

Clue from the noble gases

The early chemists found the answer when the noble gases He, Ne, Ar, Kr and Xe were discovered. Helium has two electrons in its outermost shell while all other noble gases have eight. Neither their atoms combine with themselves nor with other atoms. The probable reason for this lack of reactivity was their stability.

Duplet and octet rule

Having two electrons (for hydrogen and helium, which have only the first shell) or eight electrons in the outermost shell means stability and hence unreactivity as well. This principle was named the duplet or octet rule. Atoms form bonds because they would like to lower their energy by completing their duplet or octet.

Examples

Definition

Coordinate covalent bond is a type of covalent bond in which the shared electron pair is donated by one atom only. This bond is formed when a species has an electron pair to donate to another species. The species which donates the electron pair is called a donor, while that which accepts it is called an acceptor. An arrowhead (→) pointing towards the acceptor represents this type of bond.

1. Hydroxonium ion (H₃O⁺)

Acids provide protons (H⁺) when dissolved in water. This proton has an empty outer shell and can accept one of the two pairs of electrons present on the oxygen atom in the water molecule. As a result a hydroxonium ion is formed. The positive charge covers the whole of the hydroxonium ion, and all three bonds of oxygen behave exactly alike.

2. Ammonia and boron trifluoride

During the reaction, an electron pair from nitrogen of ammonia fills the partially empty outer shell of boron present in boron trifluoride: H₃N→BF₃.

3. Ammonium chloride

A coordinate covalent bond in ammonium chloride links nitrogen of ammonia and the proton. The positive charge is spread all over the ammonium ion. All the four bonds between nitrogen and hydrogen behave exactly alike.

Covalent vs coordinate covalent bond

This proves that the difference between a covalent bond and a coordinate covalent bond lies in the way they are formed. Once such bonds are formed, there does not remain any difference. When a coordinate covalent bond breaks, both the bonded electrons go to the donor species.

Intermolecular forces

The forces of attraction which are present between the molecules of elements and compounds are named intermolecular forces of attraction. These forces are generally very weak compared with the bonding forces between atoms. Among the three states of matter, they are weakest among the molecules of gases and strongest among the molecules of solids. The melting and boiling points of substances depend on their strength: the stronger the forces, the higher the boiling point.

1. Dipole-dipole forces

These attractive forces are present between the molecules of a polar compound like HCl. Hydrogen and chlorine attract the shared pair with different force (electronegativity). Since chlorine is more electronegative, it attracts the shared pair more strongly and the bond becomes polar: H(δ+)–Cl(δ−). Due to these partial charges the molecules of HCl start attracting each other. These forces are called dipole-dipole forces. Compounds with these forces show relatively higher melting and boiling points.

2. Hydrogen bonding

Hydrogen bonding is a special case of dipole-dipole attractive forces. When hydrogen is covalently bonded to highly electronegative elements like F, O or N, the large difference of electronegativity makes the covalent bond highly polar. As a result strong dipole-dipole attractions are observed among the molecules. In H₂O the O–H bonds are highly polar, so strong attractive forces develop between water molecules. This attractive force is called hydrogen bonding.

The strength of the hydrogen bonds causes water to have relatively higher melting and boiling points compared with H₂S and NH₃.

Non-polar covalent bond

When a covalent bond is formed between two identical atoms, the shared pair of electrons lies at the middle of the two bonded atoms. This gives rise to a pure covalent bond which is essentially non-polar in character. Covalent bonds in hydrogen and chlorine molecules are non-polar.

Polar covalent bond

When a covalent bond is formed between two non-identical atoms, the shared pair of electrons bends towards the atom which has more electronegativity. As a result a partial negative charge is created on the more electronegative atom and an equal partial positive charge on the less electronegative atom. A polar covalent bond is thus formed.

Examples

Lattice and melting point

In ionic compounds the oppositely charged ions are held together by the strong electrostatic force of attraction in the form of a crystal lattice. Since the ions are rigid, such compounds exist as very stable solids with significantly high melting points. Ions surround each other from all sides, so ionic bonds are non-directional.

Brittleness

If an external force is applied on the crystal lattice, it breaks easily: ionic solids are highly brittle.

Conduction

In the solid form ionic compounds do not conduct electricity because the ions are tightly held and cannot move. In the molten state the ions get free and start conducting electricity. In aqueous solution they also conduct, because the free ions move towards their respective electrodes.

Solubility and hydration

Ionic solids are generally soluble in water. Water not only breaks the electrostatic force of attraction but also hydrates the resulting free ions. In the process of hydration, water molecules surround and interact with ions.

Reactions in solution

Ionic compounds generally react in aqueous solution: the positive ions of one compound may react with the negative ions of the other to form a new compound. A white precipitate of silver chloride forms from NaCl and AgNO₃ solutions.

The mutually shared electrons may be shown by a dot or a cross. After mutually sharing their electrons, the bonded atoms acquire the electronic configurations of the nearest noble gas.