Ionic and covalent bonding in Solids PDF

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Table of Contents Sr. No.

Topics

Page No.

1.

Solids

03

2.

Chemical bonding of solids

03

3.

Classes of solids

03

4.

Ionic solids

03-07

5.

Properties of ionic solids

04-05

6.

Lattice energy

05-06

7.

Examples of ionic solids

06-07

8.

Applications of ionic solids

9.

Covalent solids

07-11

10.

General Properties of covalent solids

08-09

11.

Structure and properties of major covalent solids

12.

Examples of covalent solids

13.

Diamond

10

14.

Graphite

10

15.

Silicon dioxide

10

16.

Applications of covalent solids

17.

Comparison between covalent solids and ionic solids

11

18.

References

12

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09 09-10

10-11

Ionic and Covalent Bonding in Solids Solids: A solid is a state of matter characterized by particles arranged such that their shape and volume are relatively stable. A solid is one of the four fundamental states of matter, along with liquids, gases, and plasma.

Chemical Bonding of solids: The constituents of a solid tend to be packed together much closer than the particles in a gas or liquid. The reason a solid has a rigid shape is that the atoms or molecules are tightly connected via chemical bonds. The bonding may produce either a regular lattice (as seen in ice, metals, and crystals) or an amorphous shape (as seen in glass or amorphous carbon).

Classes of Solids The different types of chemical bonds that join the particles in solids exert characteristic forces that can be used to classify solids. •

Ionic Bonding: Ionic bonding is a type of chemical bonding that involves the electrostatic attraction between oppositely charged ions and is primarily interaction occurring in ionic compounds.

•

Covalent bonding: Covalent bonds (e.g., in sugar or sucrose) involve the sharing of valence electrons.

•

Metallic bonding: Electrons in metals seem to flow because of metallic bonding.

•

Van der Waals forces: Organic compounds often contain covalent bonds and interactions between separate portions of the molecule due to van der Waals forces.

IONIC SOLIDS Crystalline solids in which the particles forming the crystal are positively and negatively charged ions are called ionic solids. These ions are held together by strong electrostatic forces of attraction. These attractive forces are called ionic bonds. Examples of Ionic Solids: There are several examples of ionic solids as these are a combination of metallic and nonmetallic ions and hence the following could be considered as good examples of ionic solids. •

Sodium chloride (NaCl)

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•

Magnesium Chloride (MgCl2)

•

Calcium Fluoride (CaF2)

•

Zinc sulphide (ZnS)

Properties of Ionic Compounds The properties of ionic solids relate to how strongly the positive and negative ions attract each other in an ionic bond. Iconic solids also exhibit the following properties: Crystals Formation: Ionic solids form crystal lattices rather than amorphous solids. At an atomic level, an ionic crystal is a regular structure, with the cation and anion alternating with each other and forming a three-dimensional structure based largely on the smaller ion evenly filling in the gaps between the larger ion. An ionic crystal consists of ions bound together by electrostatic attraction. The arrangement of ions in a regular, geometric structure is called a crystal lattice. Examples of such crystals are the alkali halides, which include: potassium fluoride (KF), potassium chloride (KCl), sodium fluoride (NaF) etc. High melting and High boiling points: High temperatures are required to overcome the attraction between the positive and negative ions in ionic solids. Therefore, a lot of energy is required to melt ionic solids or cause them to boil. They have high melting and boiling points, so they are in the solid state at room temperature e.g. NaCl, 801°C and MgO2 , 852°C. High Enthalpies of fusion and vaporization: The enthalpy of fusion is the heat required melt a single mole of a solid under constant pressure. The enthalpy of vaporization is the heat required for vaporize one mole of a liquid compound under constant pressure. Just as ionic compounds have high melting and boiling points, they usually have enthalpies of fusion and vaporization that can be 10 to 100 times higher than those of most molecular compounds. They are Hard and Brittle: Ionic solids are generally hard and brittle. Both of these properties reflect the strength of the columbic force. Hardness measures resistance to deformation. Because the ions are tightly bound to their oppositely-charged neighbours and, a mechanical force exerted on one part of the solid is resisted by the electrostatic forces operating over an extended volume of the crystal.

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By applying sufficient force, one layer of ions can be made to slip over another; this is the origin of brittleness. Reason for brittleness/hardness: The slippage quickly propagates along a plane of the crystal (more readily in some directions than in others), weakening their attraction and leading to physical cleavage. Because the "ions" in ionic solids lack mobility, the solids themselves are electrical insulators. Also crystals are hard because the positive and negative ions are strongly attracted to each other and difficult to separate, however, when pressure is applied to an ionic crystal then ions of like charge may be forced closer to each other. The electrostatic repulsion can be enough to split the crystal, which is why ionic solids are brittle. Conduction of electricity in solution: An ionic solid can conduct electricity when it has melted to form a liquid, or it has dissolved in water to form an aqueous solution. Both these processes allow ions to move from place to place. When ionic compounds are dissolved in water the dissociated ions are free to conduct electric charge through the solution. Molten ionic compounds also conduct electricity. In solid state, they don’t conduct because they are strongly packed. Solubility: The solubility of these ionic solids is solvent specific as they are more readily soluble in solvents with higher dielectric constant which helps these solids to overcome high attractive force. Polar solvents with high dielectric constants are ideal whereas organic solvents with low dielectric constant are not preferred.

Lattice Energy Lattice energy is an estimate of the bond strength in ionic compounds. It is defined as the heat of formation for ions of opposite charge in the gas phase to combine into an ionic solid. For example: The lattice energy of sodium chloride, NaCl, is the energy released when gaseous Na+ and Cl– ions come together to form a lattice of alternating ions in the NaCl crystal. Na+ (g) + Cl- (g) → NaCl (s)

∆H = -787.3 kJ/mol

The negative sign of the energy is indicative of an exothermic reaction. Alternatively, lattice energy can be thought of as the energy required to separate a mole of an ionic solid into the gaseous form of its ions (that is, the reverse of the reaction shown above).

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Factors affecting Lattice energy: There are two main factors that contribute to the lattice energy of an ionic solid: ➢ Charge on the ions, ➢ Radius, or size of the ions The effect of those factors is: •

as the charge of the ions increases, the lattice energy increases

•

as the size of the ions increases, the lattice energy decreases

Lattice energies are also important in predicting the solubility of ionic solids in H 2O. Ionic compounds with smaller lattice energies tend to be more soluble in H2O. Substance

U (kJ/mol)

NaI

682

CaI2

1971

MgI2

2293

NaOH

887

Na2O

2481

NaNO3

755

Table: Representative Calculated Lattice energies of different ionic compounds

Examples of ionic solids There are several examples of ionic solids as these are a combination of metallic and nonmetallic ions and hence the following could be considered as good examples of ionic solids. Sodium Chloride (rock-salt): The most well known ionic solid is sodium chloride, also known by its geological names as rock-salt or halite. Salt (NaCl) contains positive sodium ions (Na+) and negative chloride ions (Cl-). The electron from the sodium atom transfers to the chlorine atom and the oppositely charged ions attract each other to form the NaCl ionic bond. Lithium Fluoride: This is the most "ionic" of the alkali halides, with the largest lattice energy and highest melting and boiling points. The small size of these ions (and consequent high charge densities) together with the large electronegativity difference between the two elements places a lot of electronic 5|P age

charge between the atoms. Even in this highly ionic solid, the electron that is "lost" by the lithium atom turns out to be closer to the Li nucleus than when it resides in the 2s shell of the neutral atom. Magnesium chloride: The oppositely charges of the magnesium and chloride ions attract each other and form ionic bonds.

Application of ionic solids: Applications of ionic solid are as follow: •

Calcium chloride crystals are used to melt ice and snow. The crystals lower the freezing point of water.

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Barium chloride is used to make firework. It produces green coloured explosions.

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Cobalt chloride test papers are used to detect moisture. They change colour when they absorb water.

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Potassium iodide tablets are given to the people who are exposed to high levels of radiation. They protect the thyroid gland from radiation.

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Electrolyte solutions of ionic solid are given to children who have lost ions. This can happen with vomiting or diarrhea.

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Lithium iodide is used in batteries. It is an excellent conductor of electricity.

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Sodium tetra borate or borax is used in detergents, cosmetics, and enamel glazes. It is also used in biochemistry to make buffer solution, has been used as a flame retardant, and an antifungal compound.

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Potassium nitrate or saltpetre has fond major uses in fertilizers. As an oxidizing agent it is used in rocket propellants, fireworks, and as a major constituent of gun powder.

COVALENT SOLIDS Covalent solids, also called network solids, are solids that are held together by covalent bonds. As such, they have localized electrons (shared between the atoms) and the atoms are 6|P age

arranged in fixed geometries. Distortion away from this geometry can only occur through a breaking of covalent sigma bonds. Examples: Examples of covalent solids include diamond and silica (SiO2).

General Properties of Covalent solids: Most covalent solids have low melting points and boiling points: Covalent compounds usually have low melting points and boiling points. An exception to this include molecules of silica and diamonds that have a high melting point. This can be attributed to their weak force of attraction between the various bonded atoms. Covalent solids tend to be soft and relatively flexible: This is largely because covalent bonds are relatively flexible and easy to break. As with many properties, there are exceptions, primarily when molecular compounds assume crystalline forms. The solid covalent compounds have soft structures like graphite. This is because of the presence of a cloud of electrons in between each layer of carbons atoms. When dissolved in water, covalent solids don't conduct electricity: Ions are needed to conduct electricity in an aqueous solution. Molecular compounds dissolve into molecules rather than dissociate into ions, so they typically do not conduct electricity very well when dissolved in water. The absence of charged ions is the main reason. An exception to this is graphite, where we see a cloud of electrons, it makes graphite a good conductor. Many covalent solids don't dissolve well in water: Covalent compounds do not possess polar characteristics as a general property. Therefore, these compounds are insoluble in water. Water molecules are not absolutely neutral and have a slight negative charge on the oxygen atom and slight positive charges on the hydrogen atoms and since covalent compounds are made up of neutral molecules or molecules with slight charges and hence are not attracted to water molecules strongly. Covalent solids usually have lower enthalpies of fusion and vaporization: The enthalpy of fusion is the amount of energy needed, at constant pressure, to melt one mole of a solid substance. The enthalpy of vaporization is the amount of energy, at constant pressure, required to vaporize one mole of a liquid. On average, it takes only 1% to 10% as much heat to change the phase of a molecular compound as it does for an ionic compound. Covalent solids tend to be more flammable than ionic compounds: Many flammable substances contain hydrogen and carbon atoms which can undergo combustion, a reaction that releases energy when the compound reacts with oxygen to

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produce carbon dioxide and water. Carbon and hydrogen have comparable electronegativies so they are found together in many molecular compounds.

Physical And Chemical Properties •

The liquid covalent compounds evaporate. This means the molecules of liquids and solids loses from their surface into the air.

•

These solids have very less affinity between their molecules.

•

Various covalent solids have their own characteristically shaped molecules. Their bonds are directed at pre-set angles.

•

Most of the covalent solids are non-polar or have very little tendency to split completely to form ions and hence never conduct electricity.

•

At normal temperature and pressure, we will find these compounds as either liquids or gases. But, there are solids as well and they have higher molecular weights.

•

The covalent solids crystals are of two types: One that has weak van der Waal force holding these together like in Iodine. These are easily fusible and volatile, the other having a large network of atoms setting up the macromolecules.

•

These solids are soluble in organic solvents like ether and benzene.

•

Covalent bonds are directional in nature. Therefore, they exhibit the phenomenon of isomerism.

•

Covalent solids majorly have a very slow rate of reactions, unlike the various ionic compounds.

Structure And Properties of Major Covalent Solids: There are several examples of covalent solids, some of the major covalent solids along with their structures and properties are given below:

Figure: The Structures of Diamond and Graphite. (a) Diamond consists of sp 3 hybridized carbon atoms, each bonded to four other carbon atoms. The tetrahedral array forms a giant network in which carbon atoms form sixmembered rings. (b) These side (left) and top (right) views of the graphite structure show the layers of fused sixmembered rings and the arrangement of atoms in alternate layers of graphite.

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Diamond: In diamond, each carbon atom is bonded tetrahedrally to four other carbon atoms. The carbon atoms are sp3-hybridized and held together by strong carbon-carbon single covalent bonds. The strength and directionality of these bonds make diamond the hardest known material.

Graphite: Graphite has a layer structure which is quite difficult to draw convincingly in three dimensions. Graphite has a hexagonal unit cell containing two layers offset so that the carbon atoms in a given layer sit over the middle of the hexagons of the layer below. Each carbon is covalently bonded to three other carbons in the same layer to form interconnected hexagonal rings. The Bonding in Graphite: Each carbon atom uses three of its electrons to form simple bonds to its three close neighbours. That leaves a fourth electron in the bonding level. These "spare" electrons in each carbon atom become delocalized over the whole of the sheet of atoms in one layer. The important thing is that the delocalized electrons are free to move anywhere within the sheet. The atoms within a sheet are held together by strong covalent bonds - stronger, in fact, than in diamond because of the additional bonding caused by the delocalized electrons.

Silicon Dioxides: Silicon dioxide has a giant covalent structure. Part of this structure is shown in the diagram - oxygen atoms are shown as red, silicon atoms shown as black. Each silicon atom is covalently bonded to four oxygen atoms. Each oxygen atom is covalently bonded to two silicon atoms. This means that, overall; the ratio is two oxygen atoms to each silicon atom, giving the formula SiO2.

Applications of Covalent Solids ➢ Both silicon carbide and boron carbide are hard ceramic materials that have numerous applications.

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➢ Silicon carbide is used in the manufacture of automotive brake and clutch parts, as an abrasive material, and as the main component in bullet proof vests and body armour. ➢ Boron carbide is probably the second hardest material after diamond, and is used to provide armour plating in armoured vehicles and military aircraft. ➢ Diamond is used as a gemstone in jewelry, ornaments and crowns. Diamonds are also used as abrasives to cut and polish other materials and gemstones. ➢ Diamonds being very hard is used in tools that cut glass and drill hard rocks. They are also used in knives which are used to perform sensitive operations like the cataract operation. Because of its heat sensitive properties, diamonds are used in highly sensitive thermometers and windows of space shuttle. ➢ Graphite in its powdered form is used as a lubricant in heavy machines. It is also used to make black paint and inks due to its dark, grey colour. It is used for making lead in pencils. Graphite bricks are used as moderators in atomic reactors.

Comparison between Ionic and Covalent Solids Ionic solids

Covalent solids

Ionic solids are non-conductors of electricity. However they conduct electricity in molten or

They are bad conductors of electricity except graphite.

solution form. Ionic solids have definite geometric shape.

They have definite shape. They have open structure due to

They are non-directional in shape.

valences of atoms directed in definite directions.

They do not exist in molecules due to their ionic

They may be called as molecules due

nature.

to their covalent nature.

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