Chemistry Pre-u Chemistry Sem 1 Chap 3

962 / 1 CHEMISTRY SEMESTER 1 CHAPTER 3 CHEMICAL BONDING

3.1 3.2 3.3 3.4

CHAPTER 3 : CHEMICAL BONDING Ionic bonding Covalent bonding Metallic bonding Intermolecular forces : Van der Waals forces and hydrogen bonding

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Topic P1 3. Chemical Bonding

4

P2 P1 P2 P1 P2 P1 P2 P1 P2 P1 P2

A

B, C

A

B, C

1b 5c 7a

3

19 b

3

19

3c, 6a

4

5a

3

3b, c

1

2

2, 5b

2

INTERACTION BETWEEN ELEMENTS

Metal and non-metal

Non - metal and non-metal

Metal and metal

• A Lewis dot symbol consists of the symbol of an element and one dot for each valence electron in an atom of the element

3.1 Ionic Bonding • The central idea of the ionic bonding model is the transfer of electrons from metal atoms to non-metal atoms to form ions that come together in a solid ionic compound, where ionic bond is formed in between oppositely charged ions by electrostatic attraction forces. • For example, in the formation of sodium fluoride, NaF Sodium atom Electronic configuration

Orbital diagrams

Lewis diagram

Na

1s

Fluorine atom

(1s22s22p63s1)

2s

2p

Na

F

(1s22s22p5)

Sodium fluoride Na+ F(1s22s22p6) (1s22s22p6)

3s

F

Na+

F-

• The interaction between sodium atom and fluorine atom occur, where sodium atom (with low ionisation energy) donates electron to fluorine (with high electron affinity) to form sodium ion, Na+, and fluoride ion, F-, respectively. Note that both ions have achieved octet arrangement of ns2np6, as it is the most stable form of an ion formed. • The oppositely charged Na+ and F- form a giant ionic crystal lattice, with very high melting point, via electrostatic attraction forces. i. Magnesium chloride, MgCl2

ii. potassium oxide, K2O

iii. Calcium sulphide, CaS

iv. aluminium oxide, Al2O3

• According to Coulumb's Law, electrostatic energy between two oppositely charged substance (A and B) is directly proportional to the charge carried by each ions yet inversely proportional to the distance between them

Qn+ × Qn− Electrosta tic Energy (or lattice energy ) ∝ r+ − r− • This relationship helps us predict trends in lattice energy and explain the effects of ionic size and charge a) Effect of ionic size : As we move down a group in the periodic table, the ionic radius increases. Therefore, the electrostatic energy between cations and anions decreases because the inter-ionic distance is greater; thus, the lattice energies of their compounds should decrease as well. This prediction is borne out by the alkali-metal halides. note the regular decrease in lattice energy down a group whether we hold the cation constant (LiF to LiI)

b) Effect of ionic charge : When we compare lithium fluoride with magnesium oxide, we find cations of about equal radii (Li+ = 76 pm and Mg2+ = 72 pm) and anions of about equal radii (F- = 133 pm and O2- = 140 pm). Thus, the only significant difference is the ionic charge: LiF contains the singly charged Li+ and F- ions, whereas MgO contains the doubly charged Mg2+ and O2- ions. The difference in their lattice energies is : ∆Hlattice of LiF = - 1050 kJ mol-1 ∆Hlattice of MgO = - 3923 kJ mol-1 This nearly fourfold increase in ∆Hlattice reflects the fourfold increase in the product of the charges (1 x 1 vs. 2 x 2) in the numerator of equation above

3. Properties of Ionic Compound. a) Melting point - Ionic compound has giant ionic crystal lattice, which are hold by strong electrostatic attraction forces by repeating of oppositely charged ions

Therefore, the melting point of ionic compounds are usually very high as a lot of energies are required to overcome the electrostatic attraction forces, and melted to form free moving ions. Therefore, ionic compounds have very high melting point

b) Conductivity of electricity - Most ionic compounds do not conduct electricity (insulator) in the solid state but do conduct it when melted or when dissolved in water. (except some super-ionic conductors, such as AgI, which have remarkable conductivity in the solid state.) Solid ionic salt consists of immobilized ions. When it melts or dissolves, however, the ions are free to move and carry an electric current c) Hardness and brittleness of ionic compound - All ionic solids are hard (does not dent), rigid (does not bend), and brittle (cracks without deforming). These properties are due to the powerful attractive forces that hold the ions in specific positions throughout the crystal. Moving the ions out of position requires overcoming these forces, so the sample resists denting and bending. If enough pressure is applied, ions of like charge are brought next to each other, and repulsive forces crack the ionic solid suddenly

3.2 Covalent bond • Studies of covalent bond, was widely developed ever since Lewis suggested that a chemical bond exist in a hydrogen gas occur by sharing en electron between two hydrogen atoms. • Electron pair that connect the 2 hydrogen atoms is called covalent bond, a bond in which two electrons are shared by two atoms, and the electron pair that bond between the two hydrogen atoms is also called as bonding pair electrons. • In a covalent bond, each electron in a shared pair is attracted to the nuclei of both atoms. This attraction holds the two atoms in H2 together and is responsible for the formation of covalent bonds in other molecules.

2. A Lewis structure is a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on individual atoms. Only valence electrons are shown in a Lewis structure a) Consider the fluorine molecule, F2. The electron configuration of F is 1s22s22p5. The 1s electrons are in inner shell which is nearest to the nucleus. For this reason they do not participate in bond formation. Thus, each F atom has seven valence electrons (2s22p5). Therefore, each fluorine atom needed one electron to achieve octet configuration (ns2np6)

F : 1s22s22p5

F : 1s22s22p5

1s22s22p6 1s22s22p6

• Oxygen atom has electronic configuration of 1s22s22p4. To achieve stable octet configuration (ns2np6), each oxygen atom need 2 electrons. Hence, when 2 oxygen atoms interact, they shared two electrons in between each other as described in diagram below.

O : 1s22s22p4 O : 1s22s22p4 1s22s22p6 1s22s22p6 • From the structure of oxygen molecule formed, each oxygen atom shared two electrons from each other to form a double bond, in order to achieve octet configuration among each oxygen atom

• Nitrogen has electronic configuration of 1s22s22p3 and required 3 electrons to achieve octet configuration (ns2np6). In this case, each nitrogen atom shared 3 electrons from each of its atom to form triple bond

N : 1s22s22p3

N : 1s22s22p3

1s22s22p6

1s22s22p6

Water, H2O H : 1s1 O : 2s22p4

Carbon dioxide, CO2 C : 2s22p2 O : 2s22p4

Ammonia, NH3 H : 1s1 N : 2s22p3

Ethene, C2H4 H : 1s1 C : 2s22p2

Hydrogen cyanide, HCN H : 1s1

N : 2s22p3

C : 2s22p2

Tetrachloromethane, CCl4 C : 2s22p2 Cl : 3s23p5

Ethanoic acid, CH3COOH H : 1s1 O : 2s22p4 C : 2s22p2

Ethyne, C2H2 H : 1s1 C : 2s22p2

a) Note that in ethene, hydrogen cyanide and ethanoic acid, all the valence electrons are used in bonding; there are no lone pairs on the carbon atoms. In fact, most of the stable molecules containing carbon do not have lone pairs on the carbon atoms. b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds. Bond length is defined as the distance between the nuclei of two covalently bonded atoms in a molecule. For a given pair of atoms, such as carbon and nitrogen, triple bonds are shorter than double bonds, which, in turn, are shorter than single bonds. The shorter multiple bonds are also more stable than single bonds. c) Covalent bond not only exist in neutral molecule, but also in some molecular ions. Table below shows a few example of molecular ions, which have covalent bonds in its molecule

Carbonate ion, CO32C : 2s22p2 O : 2s22p4

Cyanide ion, CNN : 2s22p3 C : 2s22p2

Sulphate ion, SO42S : 3s23p4 O : 2s22p4

Nitrate ion, NO3N : 2s22p3 O : 2s22p4

3.2.1 Exception of Octet Rules 1. From the Lewis structure sketched for sulphate ion, SO42-, we can see that the center metal atom (sulphur) has more than 8 electrons. For this, the molecule is described as molecule that can expanded octet. These molecules that have more than 8 electrons are located at Period 3 and below, as these center atoms have empty d-orbital available to expand the number of electrons positioned in the center atom. Those center atom from Period 2, such as C, N, O and F can only allocate a maximum of 8 electrons, since they do not have d-orbital available. Example : Both phosphorous and nitrogen are elements from Group 15, with the valence electron of ns2np3. They can both form NCl3 and PCl3 respectively when react with limited amount of chlorine, however under excess chlorine, only PCl5 can be formed but not NCl5. Explain the statement bolded. Solution : This is due to, phosphorous, which is from period

3, have empty d-orbital to expand the octet. However, nitrogen, which is from Period 2, do not have empty orbital and can only allocate 8 electrons in it's shell.

Phosphorous pentachloride, PCl5 P : 3s23p3 Cl : 3s23p5

Sulphur hexafluoride, SF6 S : 3s23p4 F : 2s22p5

Bromine pentachloride, BrCl5 Cl : 3s23p5 Br : 4s24p5

Xenon tetrafluoride, XeF4 Xe : 5s25p6 F : 2s22p5

2. There are also some stable covalent compounds, which have less than 8 electrons in their center atom (incomplete octet). The center atoms are usually metals with great number of valence electrons with small atomic radius, such as beryllium, boron and aluminium Be : 2s2

BeCl2 Cl : 3s23p5

BF3 B : 2s22p1

AlCl3 F : 2s22p5 Al : 3s23p1

F : 3s23p5

These compound possessed stability due to their short bond length between center atom with surround atom. Furthermore they can form resonance structure between the center atom and surrounding atoms

3. There are also some molecules which contain an odd number of electrons. Among the most common compounds are nitrogen monoxide (NO) and nitrogen dioxide (NO2). Nitrogen monoxide NO N : 2s22p3 O : 2s22p4

Nitrogen dioxide, NO2 N : 2s22p3 O : 2s22p4

Odd-electron molecules are sometimes called radicals. Many radicals are highly reactive. The reason is that there is a tendency for the unpaired electron to form a covalent bond with an unpaired electron on another molecule. For example, when two nitrogen dioxide molecules collide, they form dinitrogen tetroxide

3.2.2

Dative bond

1. As shown in the Lewis structure of nitrate ion, one of the N-O bond is drawn as →. The bond '→' placed is called as dative bond (also known as coordinative bond), where dative bond is defined as a covalent bond in which one of the atoms donate the lone pair electrons available. Although the properties of a coordinate covalent bond do not differ from those of a normal covalent bond because all electrons are alike no matter what their source. 2. Dative bond is usually applied for these few circumstances below a) To assist atom / molecule / ion that not yet achieved octet configuration. Making use of atom which has lone pair electrons to those which are lack of electrons.

Sulphur dioxide, SO2

Sulphur trioxide, SO3

Carbon monoxide, CO

O S O

O

Ozone molecule, O3

Water with hydrogen ion

Ammonia with hydrogen ion

Ammonia with boron trifluoride

b) Formation of dimer - In order for some compounds which have incomplete octet to achieve stability, they tend to form dimer or polymer among themselves by using dative bond. Two of the most common examples are aluminium trichloride and beryllium dichloride Monomer of AlCl3

Dimer of aluminium chloride, Al2Cl6

Monomer of BeCl2

Dimer of Be2Cl4

Polymer of (BeCl2)n

c) Formation of coordination compounds - Coordination compounds are substances that contain at least one complex ion, a species consisting of a central metal cation (either a transition metal or a main-group metal) that is bonded to molecules and/or anions called ligands via dative (coordinative) bond. hexaaquacopper (II) ion ; [Cu(H2O)6]2+

tetraamminenickel (II) ion ; [Ni(NH3)4]2+

Hexacyanoferrate (III) ion ; [Fe(CN)6]3-

Trioxalatocobaltate(III)ion, [Co(C2O4)3]3-

3.2.3 Hybridisation Theory Valence bond theory • The basic principle of valence bond theory is that a covalent bond is formed when orbitals of two atoms overlap and the overlapped region, which is between the nuclei, is occupied by a pair of electrons. The central themes of valence bond theory derive from this principle : – Opposing spins of the electron pair ~ Stated in Pauli's Exclusion principle, the space formed by the overlapping orbitals has a maximum capacity of two electrons that must have opposite spins. For example, when a covalent bond is formed in molecule of hydrogen, H2, the two 1s electrons of two H atoms occupy the overlapping 1s orbitals and have opposite spins

• Maximum overlap of bonding orbitals ~ The bond strength depends on attraction of the nuclei for the shared electrons, so the greater the orbital overlap, the stronger (more stable) the bond. The extent of overlap depends on the shapes and directions of the orbitals. An s orbital is spherical, but p and d orbitals have more electron density in one direction than in another. Thus, whenever possible, a bond involving p or d orbitals will be oriented in the direction that maximizes overlap. For example, in hydrogen fluoride (HF) bond, the 1s orbital of H overlaps the half-filled 2p orbital of F along the long axis of that orbital. Any other direction would result in less overlap and, thus, a weaker bond

• Hybridisation of atomic orbitals ~ To account for the bonding in simple diatomic molecules like HF, we picture the direct overlap of s and p orbitals of isolated atoms. But how can we account for the shapes of so many molecules and polyatomic ions through the overlap of spherical s orbitals, dumbbell-shaped p orbitals, and cloverleaf-shaped d orbitals? Linus Pauling proposed that the valence atomic orbitals in the molecule are different from those in the isolated atoms. The spatial orientations of these new orbitals lead to more stable bonds and are consistent with observed molecular shapes. The process of orbital mixing is called hybridisation, and the new atomic orbitals are called hybrid orbitals. Two key points about the number and type of hybrid orbitals are that i. The number of hybrid orbitals obtained equals the number of atomic orbitals mixed. ii. The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed

3.2.3.1 Type of hybridisation 1. sp3 hybridisation ~ When four electron groups surround the central atom, the center atom involved must prepare four orbitals with equal energies to overlap with the four surrounding electron groups. Valence Bond theory uses hypothetical hybrid orbitals, which are atomic orbitals obtained when two or more non-equivalent orbitals of the same atom combine in preparation for covalent bond formation. Hybridisation is the term applied to the mixing of atomic orbitals in an atom (usually a central atom) to generate a set of hybrid orbitals. In the case of sp3, We can generate four equivalent hybrid orbitals from the center atom by mixing the s orbital and the three p orbitals.

Explanation

Energy level diagram

Diagram of orbitals

a) Ground state : Carbon which act as the center atom has the valence electron of 2s22p2. b) Excited state : One of the electron from 2s is promote to 2p orbitals equal energy level c) Hybridised state : One orbital of 2s and three orbitals of 2p combined (hybrid) and rearrange themselves to the shape and orientation of a tetrahedral shape.

Molecular shape : Tetrahedral Angle between bond pair : 109.50

2. sp2 hybridisation ~ When three electron groups surround the central atom, the center atom involved must prepare three orbitals with equal energies to overlap with the three surrounding electron groups. In sp2 hybridisation, three equivalent (in terms of energy level) from the center atom by mixing the one s orbital and the two p orbitals. Using boron trifluoride (BF3) as example, sp2 hybridisation is explained

Explanation

Energy level diagram

Diagram of orbitals

a) Ground state : Boron which act as the center atom has the valence electron of 2s22p1. b) Excited state : One of the electron from 2s is promote to 2p orbitals - equal energy level

c) Hybridised state : One orbital of 2s and two orbitals of 2p combined (hybrid) and rearrange themselves to the shape and orientation of a trigonal planar shape. Molecular shape : Trigonal planar Angle between bond pair : 1200

3. sp hybridisation ~ When two electron groups surround the central atom, we observe a linear shape, which means that the bonding orbitals must have a linear orientation. VB theory explains this by proposing that mixing two nonequivalent orbitals of a central atom, one s and one p, gives rise to two equivalent sp hybrid orbitals that lie 1800 apart.

Explanation

Energy level diagram

Diagram of orbitals

a) Ground state : Beryllium which act as the center atom has the valence electron of 2s2. b) Excited state : One of the electron from 2s is promote to 2p orbitals - equal energy level

c) Hybridised state : One orbital of 2s and one orbitals of 2p combined (hybrid) and rearrange themselves to the shape and orientation of a linear shape. Molecular shape : Linear Angle between bond pair : 1800

4. The concept of hybridisation is also useful to explain molecules with double/triple bonds. By using the concept of the direct overlapping orbitals and side-touch lapping orbitals, the formation of multiple bonds in ethene, C2H4 and ethyne, C2H2, are described. a) Ethene C2H4. - Hybridisation take place for both carbon atoms in ethene molecule is sp2 hybridisation • When both hybridised C is bonded together, one of the hybridised orbital overlapped directly between each other, while the other two hybridised orbitals overlapped directly with hydrogen atoms respectively. Meanwhile. the unhybridised pz orbitals of both carbon atoms form a side-touch bond between each other, and form another bond, as shown in the diagram below. • From the diagram with C=C, there are two types of bond. A sigmabond (σ-bond) is covalent bonds formed by orbitals overlapping endto-end, with the electron density concentrated between the nuclei of the bonding atoms, while the second type is called a pi bond (πbond), which is defined as a covalent bond formed by sideways overlapping orbitals with electron density concentrated above and below the plane of the nuclei of the bonding atoms

Explanation

Energy level diagram

Diagram of orbitals

a) Ground state : Carbon which act as the center atom has the valence electron of 2s22p2. b) Excited state : One of the electron from 2s is promote to 2p orbitals - equal energy level

c) Hybridised state : One orbital of 2s and two orbitals of 2p combined (hybrid) and rearrange themselves to the shape and orientation of a trigonal planar shape. Note that on unhybridised pz orbital, an electron is presence. Molecular shape : Trigonal planar Angle between bond pair : 1200

b) Ethyne, C2H2, has the Lewis structure of H−C≡C−H, which the bonding can be explain using sp hybridisation. Table below described how the hybridisation take place on each carbon atom and how the formation of triple bond occur. c) From the diagram, formation of -C≡C- is due to the formation of one sigma-bond, by direct overlapping of one of the two hybrid orbitals on each C, while the other two pi-bond bonds are formed as a result of side-lapping of the each two unhybridised py and pz orbitals

Explanation

Energy level diagram

Diagram of orbitals

a) Ground state : Carbon which act as the center atom has the valence electron of 2s22p2. b) Excited state : One of the electron from 2s is promote to 2p orbitals - equal energy level c) Hybridised state : One orbital from 2s and 2p orbitals combined (hybrid) and rearrange themselves to the shape and orientation of a linear shape. Note that on unhybridised py & pz orbitals, electrons are presence to form two π-bonds.

py

pz

sp

Molecular shape : Linear Angle between bond pair : 1800

5. Other examples of applications in valence bond theory includes the formation of nitrogen molecule, N2 and hydrogen cyanide, HCN molecule. a) Nitrogen gas is Earth's most abundant gas as it cover 78% of the content of our air. Nitrogen molecule is an inert gas thanks to its short covalent bond and also its strong triple bond. Therefore, a lot of heats are required to break the chemical bond of nitrogen before it can be applied in industries. Using valence bond theory, the bonding of nitrogen molecule is explained in the diagram below b) So, when the two hybridised nitrogen atom interacting among each other, they hence formed a linear shape, and two pi-bonds (π−bonds) are formed as a result of side-lapping of unhybridised py and pz orbital respectively

N

N

Explanation

Energy level diagram

Diagram of orbitals

a) Ground state : Both nitrogen which act as the center atom has the valence electron of 2s22p3 b) Excited state : One of the electron from 2s is promote to 2p orbitals - equal energy level c) Hybridised state : One orbital of 2s and one orbitals of 2p combined (hybrid) and rearrange themselves to the shape and orientation of a linear shape.

Molecular shape : Linear Angle between bond pair : 1800

b) Whereas for hydrogen cyanide, HCN, both carbon atom and nitrogen atom undergoes sp hybridisation, in order to form a linear structure. Explanation a) Ground state : Valence electron C : 2s22p2 N : 2s22p3 b) Excited state : One of the electron from 2s is promote to 2p orbitals - equal energy level c) Hybridised state : One orbital of 2s and one orbitals of 2p combined (hybrid) and form sp hybridisation

Energy level diagram for carbon atom, C

Energy level diagram for nitrogen atom, N

π σ

π σ

N

6. However, there are a few limitation on valence bond theory, such as when explaining the effect angle of bond-pair and bond-pair electrons when there is/are presence of lone pair electron in the center atom, and also the difference of electronegativity. a) Bonding in ammonia, NH3 and water, H2O. ~ Both NH3 and H2O undergoes an arrangement similar to sp3 hybridisation, similar to that of C in methane. Table below shows the hybridisation occur for nitrogen in ammonia, and oxygen in water Explanation a) Ground state : Valence electron N : 2s22p3 O : 2s22p4 b) Excited state : One of the electron from 2s is promote to 2p orbitals

Energy level diagram for nitrogen atom, N

Energy level diagram for oxygen atom, O

c) Hybridised state : One orbital of 2s and three orbitals from 2p combined (hybrid) and rearrange themselves to the shape and Similar to arrangement in orientation similar to that tetrahedral of tetrahedral shape.

Similar to arrangement in tetrahedral

Diagram of overlapping of ammonia, NH3 and water, H2O

Shape and angle Number of bond pair & lone pair electrons

Shape : trigonal pyramidal Angle : 1070

Shape : bent Angle : 104.50

Bond pair electron : 3 Lone pair electron : 1

Bond pair electron : 2 Lone pair electron : 2

• The angle of bond pair - bond pair electrons in ammonia and water are 1070 and 104.50 respectively, which is lesser than in methane molecule (109.50). This can be explained by the fact of the presence of lone-pair electrons in both ammonia and water. Since the lone pair - lone pair electron repulsion is stronger than lone pair - bond pair electron repulsion than bond pair - bond pair electron repulsion, it is expected that the repulsion between lone pair - bond pair electrons in ammonia is stronger than bond pair - bond pair electrons repulsion, hence caused the angle to "squeeze" to a smaller angle. As for water, since there is a presence of lone pair - lone pair electrons repulsion, it results the bond pair - bond pair electron repulsion to be much smaller, hence caused the angle to "squeeze" to a smaller angle. Though, valence bond theory does not actually explain the hybridisation especially when it relates to the repulsion occur involving lone pair electron, another theories shall be applied to study such effect. All these shall be discussed further in VSEPR theory

b) Effect of electronegativity and bonding angle between bond pair - bond pair electrons ~ i. As discussed earlier, in the bonding of water (H2O), the molecular shape and angle is described in the diagram below.

Sulphur, which is also an element from Group 16, formed hydrogen sulphide, H2S when sulphur react with hydrogen. However, unlike water, the bonding angles is much smaller compare to water. This is due to the difference of electronegativity and also the bond length between O and S in molecule. Since O is more electronegative compare to S, the O-H bond is pulled closer toward O. Furthermore, the bond length of O-H is shorter compare to S-H. As a result, the bonding pair-bonding pair electrons repulse greater with each other in H-O-H, and caused the angle become greater

ii. Another example is between NH3 and PH3. The orbital diagram for both NH3 and PH3 are described below. Both nitrogen, N, and phosphorous, P, are from the same group, which is Group 15.

Similar to the above case, P in phosphine is less electronegative than N in ammonia, and bond length of N-H is shorter than P-H. As a result, H is pull closer to N and repulsion between bonding of H-N-H is greater, compare to H-P-H, and caused the angle between H-N-H is greater compare to H-P-H

3.2.4 Valence Shell Electron Pair Repulsion (VSEPR) Theory 1. Molecular geometry is the three-dimensional arrangement of atoms in a molecule. A molecule’s geometry affects its physical and chemical properties, such as melting point, boiling point, density, and the types of reactions it undergoes. The basic concept of VSEPR theory is based on the three general rules below a) As far as electron-pair repulsion is concerned, double bonds and triple bonds can be treated like single bonds. This approximation is good for qualitative purposes. However, you should realize that in reality multiple bonds are “larger” than single bonds; that is, because there are two or three bonds between two atoms, the electron density occupies more space. b) If a molecule has two or more resonance structures, we can apply the VSEPR model to any one of them. Formal charges are usually not shown. c) The order of repulsion strength of lone pair and bond pair are : lone-pair & lone-pair electrons repulsion are the strongest, followed by lone-pair & bond-pair electrons repulsion, while bond-pair & bond-pair electrons repulsion is the weakest.

2. In order to deduce the geometry of molecule or ion, we should know the number of electrons surrounding a central atom in its Lewis structure. Step 1 : Calculate the total number of valence electrons from all atoms in a molecule. If it is a molecular ion, depend on the type of ions, the number of electrons are added / subtracted. If the molecular ion is positively charged, the total number of electrons are subtracted, however, if its negatively charged, the total number of electrons are added. Step 2 : Arrange all the atoms surrounding the central atom by using a pair of electron per bond (- 2 electrons per bond formed) Step 3 : Assign the remaining electrons to the terminal atoms so that each terminal atom has 8 electrons [except for hydrogen] (- 6 electrons per surround atom) Step 4(a) : Place any left-over electron on the central atom. (or) Step 4(b) : Form multiple bonds if there are not enough electrons to give the central atom an octet of electrons

Class

AB2

AB3

AB2E

AB4

No of No of lone surroun pair d electro atoms n 2

3

2

4

Molecular geometry

Diagram of the molecular shape

Example of molecules

0

Linear

CO2 BeCl2

0

Trigonal planar

AlCl3 BF3 NO3-

1

Shape : Bent

SO2 O3 NO2-

0

Shape : Tetrahedral

CH4 SiCl4 SO42-

AB3E

AB2E2

3

2

Shape : Trigonal pyramidal

NH3 PCl3 SO32-

2

Shape : Bent

H2O SCl2 H2O2

PCl5 SbF5

SCl4 PF4-

1

AB5

5

0

Shape : Trigonal bipyramidal

AB4E

4

1

Shape : See-Saw

AB3E2

AB2E3

3

2

2

Arrangeme nt: Trigonal bipyramidal Shape : T-shape

ICl3 BrF3

3

Arrangeme nt: Trigonal bipyramidal Shape : linear

I3BrCl2-

AB6

AB5E

AB4E2

6

5

4

0

Arrangeme nt & Shape : Octahedral

SF6

1

Arrangeme nt: Octahedral Shape : Square pyramidal

SbCl52IF5

2

Arrangeme nt: Octahedral Shape : Square planar

XeF4 BrF4-

a) phosphorous trichloride, PCl3 S1. Total valence electrons P = 5 e- ; 3 Cl = 3 x 7eTotal electrons = 26 S2. Electrons used / bond = 3 x 2eElectrons left = 26 - 6 = 20 e-

S3. e- at surround atom = 3 x 6eElectrons left = 20 - 18 = 2 e-

S4a ) The 2e- remain is placed at the center atom P

Since molecule contain 3 surrounding atom and 1 lone pair electrons, hence

Arrangement : tetrahedral Shape : trigonal pyramidal

b) Carbonate ion, CO32-. S1. Total valence electrons C = 4 e- ; 3 O = 3 x 6e- + 2e- accept ; Total electrons = 24

S2. Electrons used / bond = 3 x 2eElectrons left = 24 - 6 = 18 e-

S3. e- at surround atom = 3 x 6eElectrons left = 18 - 18 = 0 e-

S4b) Since the center atom C not yet achieved octet, a double bond is form using any e- from O Since molecular ion contain 3 surrounding atom and 0 lone pair electrons, hence

Arrangement and Shape : trigonal planar

c) iodine tetrachloride ion, ICl4S1. Total valence electrons I = 7 e- ; 4 Cl = 4 x 7e- ; + 1e- accept ; Total electrons = 36

S2. Electrons used / bond = 4 x 2eElectrons left = 36 - 8 = 28 e-

S3. e- at surround atom = 4 x 6eElectrons left = 28 - 24 = 4 e-

S4a ) The 2e- remain is placed at the center atom P

Since molecular ion contain 4 surrounding atom and 2 lone pair electrons, hence

Arrangement : octahedral Shape : square planar

a) Iodide ion, I3-

b) Antimony pentachloride ion, [SbCl5]2-

Arrangement : trigonal bipyramidal Shape : linear c) Bromine trifluoride, BrF3

Arrangement : trigonal bipyramidal Shape : T-shape

Arrangement : octahedral Shape : square pyramidal d)

sulphur tetrafluoride, SF4

Arrangement : trigonal bipyramidal Shape : see-saw

3.2.5 Electronegativity and Polarity of Molecules 1. From all the chemical bonding discussed so far, ionic and covalent bonding models portray compounds as being formed by either complete electron transfer or complete electron sharing. However, in most real compounds, the type of bonding lies somewhere between these extremes. Thus, the great majority of bonds are more accurately thought of as “polar covalent,” that is, partially ionic and partially covalent. Pure ionic compound Pure covalent compound

Polar covalent compound

2. One of the most important concepts in chemical bonding is electronegativity (EN), the relative ability of a bonded atom to attract the shared electrons. Electronegativity is a relative concept, meaning that an element’s electronegativity can be measured only in relation to the electronegativity of other elements. Linus Pauling devised a method for calculating relative electronegativities of most elements

Molecule

Fluorine, F2

δ+

Lewis structure

Polarity ∆EN

Hydrogen fluoride, HF δ−

Non-polar molecule

Polar molecule

4.0 - 4.0 = 0

4.0 - 2.1 = 1.9

Since there are no different between the EN, therefore bonding pair Explanatio electrons was not pulled to n either atom, hence remain in the middle between 2 F atom Dipole Since there is no moment, difference between EN, magnitude the dipole moment is 0, and and no resultant dipole vector. moment nor vector

Since F is more electronegative than H, therefore, the bonding pair electrons were pulled closer to F atom. This will caused F to have greater electron density compare to H. Therefore, F has partial negative charge (δ−) while H carries partial positive charge (δ+) Since F is more electronegative than H. There is presence of dipole moment in HF and the vector of resultant dipole moment is pointed to the direction of F (symbolised by I ).

a) Comparisons above are basically the difference between an element with compound, where diatomic molecules containing atoms of different elements (for example, HCl, CO, and NO) have dipole moments and are called polar molecules while diatomic molecules containing atoms of the same element (for example, H2, O2, and F2) are examples of non-polar molecules because they do not have dipole moments. However, not necessarily, a covalent bond compound is guaranteed a polar molecule. b) For a molecule made up of three or more atoms both the polarity of the bonds and the molecular geometry determine whether there is a dipole moment. Even if polar bonds are present, the molecule will not necessarily have a dipole moment. For example, comparison between sulphur dioxide and sulphur trioxide

Molecules

Sulphur trioxide, SO3

Sulphur dioxide, SO2

Non-polar molecule S = 2.5 ; O = 3.5 the dipole moment of the entire molecule is made up of three bond moments, that is, individual dipole moments in the polar S−O bonds. The bond moment is a vector quantity, which means that it has both magnitude and direction. The measured dipole moment is equal to the vector sum of the bond moments. The three bond moments in SO3 are equal in magnitude. Because they point in opposite directions in a planar SO3 molecule, the sum of resultant dipole moment would be zero

Polar molecule S = 2.5 ; O = 3.5

Lewis structure and shape Polarity EN

Bond moment, dipole moment magnitude and vector

the dipole moment of the entire molecule is made up of two bond moments and one lone pair electron. Even though two bond moments in SO2 are equal in magnitude, however, the presence of the lone pair electrons which caused the repulsion of bond pair electrons to be lesser. Because they point in downward directions in a bent SO2 molecule, the overall vector points downward and the sum of resultant dipole moment would not be zero, hence a polar molecule

From the example of SO2 and SO3 used, we can tell that if a polyatomic molecule is a symmetrical molecule (molecule with no lone pair electrons in it), it may be a non-polar molecule. However if a polyatomic molecule is an asymmetrical molecule (molecule with lone pair electrons in it), it may be a polar molecules. c) Even though, a polyatomic molecule may be symmetrical, if the surrounding atoms are not the same, molecule may be a polar molecule, as the bonding moments are different, and caused the magnitude of dipole moment of the molecule is not equal to zero. However, if the surrounding atoms are the same, bonding moments are equal in magnitude, and the resultant vector cancel-off each other, causing the dipole moment is equal to zero, hence form a non-polar molecule. For example

Molecule

Methane, CH4

Chloromethane (chloroform), CH3Cl

Non - polar molecule

polar molecule

Lewis structure

Polarity

As methane is a symmetrical molecule, and the surrounding atoms are the same, the vector of bond Explanation moment cancel off each other, hence caused the dipole moment is equal to zero.

Since a foreign element, Cl, is in the symmetrical molecule, and Cl is more EN than the rest of the atoms, the vector and magnitude is heading to the direction of Cl, caused a small dipole moment present in molecule, hence polar.

Covalent molecule

Diatomic molecule

Same element

Non-polar molecule

Polyatomic molecule

Different element

Asymmetric al

Polar molecule

Polar molecule

Symmetrical

Same surround atoms

Different surround atoms

Non-polar molecule

Polar molecule

3.2.6.

Electronegativity and Type of Chemical Bond.

1. The type of bond that would form can be told by using the difference of electronegativity (∆EN). Larger the difference, the more tendency of electron from low electronegativity atom to move to the atom with higher electronegativity and form ionic compound. a) The relationship between the ionic character and the difference in the electronegativity of the bonded atom is shown in the diagram and graph below.

b) From the graph above, the dotted line represent the arbitrary line between ionic and covalent characteristic of a molecule. To be more specific, there more likely an ionic compound may have high covalent characteristic (exemplified by LiI), or conversely covalent compound having high ionic characteristic (exemplified by HF). c) The covalent characteristic of a molecule is dependent on the ability of a cation to polarise an anion. Polarisation indicates the ability of a cation to attract the electron density of an anion when put next to the cation involved. When a cation is able to pull the electron density of the anion closer to it, as if the anion wanted to share electron with cation, hence increase the covalency of the molecule

A+

X–

B+

Y-

The covalency properties of a molecule is dependent on the cation and anion where they can be explained qualitatively via • Polarisation power of cation • Polarisability of anion

3.6.1.1 Polarisation Power of Cation Polarisation Power of Cation – measure the ability of a cation to polarise the electron cloud of the anion. 2 factors determining the polarisation power of cation Charge of cation Size of cation ⇒ Greater the charge of ion, higher the ⇒ Smaller the size of cation, closer the effective nuclear charge of cation, neighboring anion to the nucleus of hence it will be able to attract the cation, hence easier for the cation to neighboring electron density of anion. polarise the anion and result an This will caused the polarization power increment in the polarization power of of cation increase, hence increase the cation, and increase the covalent covalent characteristic of cation. characteristic of cation. ♦ Both factors can be explained in another term called as charge density where Charge Density = Charge / Ionic Radius ♦ From the equation above, Charge Density will have a greater value, provided that cation has a high charge and small cationic radius. ♦ Greater the charge density, higher the polarization power, greater the covalent characteristic of the cation.

3.6.1.2 Polarisability of Anion • Polarisability of an anion ~ ability of the anion to allow the electron density to be polarised by cation. • 2 factors determining the polarisability of an anion Charge of anion

Size of anion

⇒ Greater the charge of anion, lower the ⇒ Larger the size of anion, further the effective nuclear charge of anion. This will outermost electron from the nucleus weakened the electrostatic attraction forces of the anion, easier for the cation to between nucleus and the outermost polarise the anion, and cause the electron in anion, and increase the polarisability to increase, hence polarisability of the anion, hence increase increase the covalent characteristic the covalent characteristic of anion of anion.

• Unlike cation, anion does not have a term that combined both factors of charge and ionic radius. However, information of polarisability of anion enable the prediction of the covalent characteristic of a molecule, since in order to form a covalent bond, it depend on both polarisation power of cation and polarisability of the anion

3.6.2 Prediction of Chemical Bond :Fajans’ Rule • In 1923, Kazimierz Fajans formulated an easy guidance to predict whether a chemical bond will be covalent or ionic, and depend on the charge on the cation and the relative sizes of the cation and anion. They can be summarized in the following table Ionic compound

Low positive charge

Large cation

Small anion

Covalent compound

High positive charge

Small cation

Large anion

• Based on these guidance, the bonding of a few compounds shall be discussed to understand the application of Fajans’ Rule in the chemical bonding

Lithium halide (LiX) • Lithium ion, Li+ (1s2) has a small size due to only 1 shell present in its ion. But since it has a low charge, so its charge density is not too high. That is why, all lithium halide are ionic compound. The covalency of lithium halide varies from a highly ioniccharacteristic to highly covalency, depending on the polarisability of the anion next to Li+ • When a group of halide, F– ; Cl–; Br–; I– is put close to Li+, the covalency of lithium halide increase when going down to Group 17 halide. LiF is highly ionic, since the fluoride ion has small ionic size and low charge, hence has low polarisability. Ionic size increase with the increasing shell when going down to Group 17 halide, hence increase the polarisability, which allowed lithium ion to polarise the anion’s electron density, hence increase the covalency

Cl– Br–

F–

Li+

Aluminium halide (AlX3) and aluminium oxide (Al2O3) • Aluminium ion (Al3+) has high charge density, due to its high charge unit and its small ionic radius. So, depending on the anion, aluminium has a high tendency to form covalent compound. For example, when going down to Group 17 halide, aluminium fluoride (AlF3) forms ionic compound (since F- has a low polarisability), while aluminium trichloride (AlCl3), aluminium tribromide (AlBr3) and aluminium iodide (AlI3) form covalent compound (since chloride, bromide and iodide have high polarisability). This explained why aluminium fluoride has a high melting point (10400C), while aluminium trichloride and tribromide are 1920C and 780C respectively. • As for aluminium oxide (Al2O3), it is an ionic compound with high covalent characteristic, as aluminium ion has high covalent characteristic due to its high charge density. This explained the high melting point of Al2O3 (20500C) yet it is insoluble in water. It also explained the amphoteric properties of aluminium oxide where aluminium oxide can act as an acid (covalent characteristic), as well as a base (ionic characteristic).

3.3 Metallic Bonding 1. Metallic bonding occurs when large numbers of metal atoms interact. Unlike the reaction of metal with non metal which involve electrons transfer, when two metal atoms interact, they can also share their valence electrons in a covalent bond and form gaseous, diatomic molecules of M2. The electron-sea model of metallic bonding proposes that all the metal atoms in the sample contribute their valence electrons to form an electron “sea” that is delocalized throughout the piece. The metal ions are submerged within this electron sea in an orderly array

a) The model we will use to study metallic bonding is band theory because it states that delocalized electrons move freely through “bands” formed by overlapping molecular orbitals. b) Consider magnesium, for example. The electron configuration of Mg is 1s22s22p63s2, so each atom has two valence electrons in the 3s orbital. In a metallic crystal, the atoms are packed closely together, so the energy levels of each magnesium atom are affected by the immediate neighbors of the atom as a result of orbital overlaps . These molecular orbitals are so closely spaced on the energy scale that they are more appropriately described as a “band”. The closely spaced filled energy levels make up the valence band. The upper half of the energy levels corresponds to the empty, delocalized molecular orbitals formed by the overlap of the 3p orbitals. This set of closely spaced empty levels is called the conduction band. As a conductor, the conduction band and valence band are overlapped, hence electrons can travel freely among the two bands, hence conduct electricity.

c) Theoretically, greater the number of valence electrons in a metal, greater the number of electrons delocalised, higher the conductivity. However, the conductivity decrease with temperature as vibration of the lattice of ion impedes the free movement of electron in conduction band.

2. Semiconductors are element that normally are not conductors, but will conduct electricity at elevated temperatures or when combined with a small amount of certain other elements. These elements are usually metalloid such as silicon and germanium. (a) The energy gap between the conduction band and valence band of these solids is much smaller than that for insulator. If the energy needed to excite electrons from the valence band into the conduction band is provided, the solid becomes a conductor. Note that this behavior is opposite that of the metals. (b) Addition of impurities (doping) to semiconductors also increases the conductivity of the semiconductors. Doping can be done by adding one of the following: i. dopant atoms containing fewer valence electrons. Hence the semiconductor formed is positive, p - type semiconductor ii. dopant atoms with extra valence electrons. Hence the semiconductor formed is negative, n - type semiconductor

3.

Insulators are substances that do not conduct electricity no matter how high the temperature is applied to the substances involved. The energy gaps between the conduction band and valence band of these element is very large, hence regardless how much energies were applied to these insulator, it will not be able to conduct electricity nor heat. Glass and woods are good examples of insulator. In wood and glass, the gap between the valence band and the conduction band is considerably greater than that in a metal. Consequently, much more energy is needed to excite an electron into the conduction band. Lacking this energy, electrons cannot move freely. Therefore, glass and wood are insulators, ineffective conductors of electricity

3.4 Intermolecular forces : Van der Waals forces and hydrogen bonding 1. The nature of the state of matter of substances and their changes are due primarily to forces among the molecules. Both bonding (intramolecular) forces and intermolecular forces arise from electrostatic attractions between opposite charges. Bonding forces are due to the attraction between cations and anions (ionic bonding), nuclei and electron pairs (covalent bonding), or metal cations and delocalized valence electrons (metallic bonding). Intermolecular forces, on the other hand, are due to the attraction between molecules as a result of partial charges, or the attraction between ions and molecules. The two types of forces differ in magnitude, and forces explains why: a) Bonding forces are relatively strong because they involve larger charges that are closer together. b) Intermolecular forces are relatively weak because they typically involve smaller charges that are farther apart

3. Induced dipole Forces (Dispersion forces) - Consider how a helium atom (monoatomic gas which have dipole moment = 0), interact with the following species. Helium with cation

He

He

Helium with polar molecule

Helium with Helium

a) From the diagram, we can tell that if an ion or a polar molecule is placed near an atom or a non-polar molecule, the electron distribution of the atom (or molecule) is distorted by the force exerted by the ion or the polar molecule, resulting in a kind of dipole. The dipole in the atom (or non-polar molecule) is said to be an induced dipole because the separation of positive and negative charges in the atom (or non-polar molecule) is due to the proximity of an ion or a polar molecule. b) However, the weak attractive interaction between a non-polar molecule to another non-polar molecule are unlike when placed near an ion or polar molecule. Between two non-polar atom, they form among themselves a short induced dipole, hence attract each other temporary. Therefore, the forces formed between them are very weak and can be broken easily. Such interaction is also known as London forces

Alkane

CH4

C2H6

C3H8

C4H10

C5H12

C6H14

C7H16

C8H18

RMM

16

30

44

58

72

86

100

114

Boiling point oC

– 162

– 8.6

– 42.2

– 0.5

36.3

68.7

98.4

126

Trend

Relative molecular mass increased weak Van Der Waals' forces increased

d) For molecules, especially organic compounds, which have the same molecular mass and functioning group, they may have different boiling point, depend on the molecular structure. For example, pentane, C5H12, with molecular mass 72, has 3 isomers, as shown in table below. Molecule

2,2-dimethylpropane

2-methylbutane

n-pentane

5.4

21.8

36.3

Molecular structure

Boiling point / oC

Total surface area increased weak Van Der Waals' forces increased Boiling point increased

4. Dipole-Dipole Forces ~ When polar molecules lie near one another, as in liquids and solids, their partial charges act as tiny electric fields that orient them and give rise to dipoledipole forces: the positive pole of one molecule attracts the negative pole of another. These are the forces that give polar compound a higher boiling point than the non-polar compound.

Molecule

Formula

Molecular mass

Dipole moment

Boiling point (K)

Propane

CH3CH2CH3

44

0.08

231

Dimethyl ether

CH3OCH3

46

1.3

248

Methyl chloride

CH3Cl

50.5

1.87

249

Ethanal

CH3CHO

44

2.69

294

Acetonitrile

CH3CN

41

3.92

355

Almost the same molecular mass However, greater the dipole moment Stronger the dipole-dipole forces Higher the boiling point

5. Hydrogen bond ~ a special type of dipole-dipole interaction between the hydrogen atom in a polar bond, as in N−H, O−H, or F−H, with an electronegative O, N, or F atom. Diagram below shows a few example of interaction between molecules using hydrogen bond Hydrogen bond

Even though F is more electronegative than O & Hydrogen bond between F–H is stronger than O–H However, water form more hydrogen bond Strong hydrogen bond

F is more electronegative than N Hydrogen bond between F–H is stronger than N–H Higher the boiling point

b) The factors of hydrogen bonding can also use to explain the solubility of some organic compound in water, like example, ethane cannot dissolve in water but ethanol can dissolve in water, due to the hydrogen bonding.

c) Some organic compound form dimer using hydrogen bond. For example, when glacial ethanoic acid is dissolved in organic solvent, it form a dimer using hydrogen bond via the interaction between O-H and C=O in each of the molecule.

d) In some case, hydrogen bond can also be used to form which is the intermolecular forces and intramolecular forces. For example, in 2-nitrophenol and 4-nitrophenol, the boiling point of the 2 compounds can be explain below :

Since 2-nitrophenol form strong hydrogen bond as intramolecular forces, the interaction between 2-nitrophenol molecules are weaker among each other, compare to 4nitrophenol, which used hydrogen bond as their intermolecular forces. With stronger hydrogen bond which act as the intermolecular forces, the boiling point of 4-nitrophenol is expected to be higher than 2-nitrophenol