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Pioneer Junior College H2 Chemistry 9647: Chemical Bonding

of 32 PJC Chemistry Unit 2012

Pioneer Junior College Chemistry Higher 2 (9647)

Chemical Bonding

References

1. Longman A-Level Guide, Chemistry – J.G.R. Briggs 2. E.N Ramsden: A-Level Chemistry 3. Graham Hill & Holman: Chemistry in Context Content

• Metallic bonding

• Ionic (electrovalent) bonding

• Covalent bonding and co-ordinate (dative covalent) bonding

(i) The shapes of simple molecules

(ii) Bond energies, bond lengths and bond polarities

• Intermolecular forces, including hydrogen bonding

• Bonding and physical properties

• The solid state

Assessment Objectives

Candidates should be able to:

(a) describe ionic (electrovalent) bonding, as in sodium chloride and magnesium oxide, including the use of ‘dot-and-cross’ diagrams

(b) describe, including the use of ‘dot-and-cross’ diagrams, (i) covalent bonding, as in hydrogen; oxygen; nitrogen; chlorine; hydrogen chloride;

carbon dioxide; methane; ethene (ii) co-ordinate (dative covalent) bonding, as in formation of the ammonium ion and

in the Al2Cl6 molecule.

(c) explain the shapes of, and bond angles in, molecules such as BF3 (trigonal planar); CO2 (linear); CH4 (tetrahedral); NH3 (trigonal pyramidal); H2O (non-linear); SF6 (octahedral) by using the Valence Shell Electron Pair Repulsion theory

(d) describe covalent bonding in terms of orbital overlap, giving σ and π bonds (e) predict the shapes of, and bond angles in, molecules analogous to those specified in (c) (f) describe hydrogen bonding, using ammonia and water as examples of molecules

containing -NH and -OH groups (g) explain the terms bond energy, bond length and bond polarity and use them to compare

the reactivities of covalent bonds (h) describe intermolecular forces (van der Waals’ forces), based on permanent and

induced dipoles, as in CHCl3(l); Br2(l) and the liquid noble gases

(i) describe metallic bonding in terms of a lattice of positive ions surrounded by mobile electrons

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

(k) deduce the type of bonding present from given information (l) show understanding of chemical reactions in terms of energy transfers associated with

the breaking and making of chemical bonds



(m) describe, in simple terms, the lattice structure of a crystalline solid which is: (i) ionic, as in sodium chloride, magnesium oxide (ii) simple molecular, as in iodine (iii) giant molecular, as in graphite; diamond (iv) hydrogen-bonded, as in ice (v) metallic, as in copper [the concept of the ‘unit cell’ is not required]

(n) outline the importance of hydrogen bonding to the physical properties of substances, including ice and water

(o) suggest from quoted physical data the type of structure and bonding present in a substance

(p) recognise that materials are a finite resource and the importance of recycling processes

I. INTRODUCTION

1.1 Why do atoms combine? Atoms combine to achieve an outer electronic configuration resembling that of a noble gas, i.e. duplet or octet structure. Octet rule: Atoms / ions with 8 electrons in the valence shell are found to be stable (i.e. ns2 np6 valence shell electronic configuration). Duplet rule: Similar stability is found in atoms / ions with 1s2 electronic configuration e.g. H2,

He, Li+. Electron rearrangement occurs in two ways:

1) of electrons from an atom to another atom 2) of electrons between two or more atoms Electron rearrangement results in the formation of chemical bonds.

1.2 What is a chemical bond? A chemical bond is an that binds together two or more atoms, ions or molecules. It essentially arises due to the attraction between opposite charges, either between electrons and nuclei, or as the result of a dipole attraction. Type of chemical bonds

Bonds between/ among particles (atoms or ions)

Metallic

Ionic / Electrovalent

Covalent

Intermolecular forces between molecules

van der Waals’ forces

o Permanent dipole – permanent dipole attraction

o Instantaneous dipole – induced dipole attraction

Hydrogen bond



1.3 Strength of chemical bonds

Type of bonds Strength of bonds Bond energy / kJ mol-1

Ionic bond

strong

100 – 450

Covalent bond 100 – 550

Metallic bond 50 – 600

Hydrogen bond

weak

13 – 30

van der Waals’ forces (permanent dipole – permanent dipole attraction)

5 – 12

van der Waals’ forces (instantaneous dipole – induced dipole attraction)

< 4

II. Metallic Bonding

(i) describe metallic bonding in terms of a lattice of positive ions surrounded by mobile electrons

(m) describe, in simple terms, the lattice structure of a crystalline solid which is: (v) metallic, as in copper

(p) recognise that materials are a finite resource and the importance of recycling processes

Metallic Bond is the electrostatic attraction between positively charged cations and the negatively charged sea of delocalised mobile electrons.

2.1 Structure of Metals Metallic structure consists of a giant lattice of metal surrounded by a sea of delocalised mobile electrons (Fig. 1). The strong electrostatic attraction between the positively charged nuclei and the delocalised mobile electrons gives rise to metallic bond.

Fig. 1 http://www.lstlcw.edu.hk/t9544/animation/metallic_bond/metallic.html

2.2 Nature of the metallic bond The metallic bond is strong and non-directional.

Factors affecting strength of metallic bond:

(a) Radius of metal ion: The the radius of the metal ion, the stronger is the attraction for the

delocalised electrons, hence the stronger the metallic bond.

(b) Number of valence electrons contributed by each atom:

The valence electrons contributed by each metal atom, the stronger is the metallic bond.

metal ions

delocalised mobile electrons

http://www.lstlcw.edu.hk/t9544/animation/metallic_bond/metallic.html



Metal Electronic Configuration No. of valence electrons per atom Melting

point / C

Li [He] 2s1 1 181 Na [Ne] 3s1 1 98 Mg [Ne] 3s2 2 650 Al [Ne] 3s2 3p1 3 660

The strength of metallic bond is reflected in the physical properties like hardness, high melting and boiling points.

2.3 Properties of Metals High melting and boiling points

This is due to the electrostatic attraction between cations and mobile electrons. Therefore, large amount of are needed to weaken the metallic bonds.

Good conductor of electricity (in the solid and molten state)

This is due to presence of as charge carriers. When a potential difference is applied across the ends of a metal, the delocalised electrons will flow towards the positive end, e.g. copper used in electrical wires and cables.

High thermal conductivity Fig. 2

When heat is supplied at one end of the metal, of the atoms and electrons is increased. These particles collide, transfer and conduct heat energy to surrounding atoms towards the cooler region of the metal. E.g. copper is used in pipes and radiators in central heating system, aluminium conducts heat well and is used as packaging foils and overhead cables.

Malleable (hammered into sheets) and ductile (drawn into wires)

If stress is applied to metals, layers slide over each other without overcoming the metallic bond.

Fig. 3 Shiny appearance

The delocalised electrons are in energy levels which are close to each other. Thus, they absorb radiation of almost all frequencies and would be excited to higher energy levels. As electrons return to lower energy levels, is emitted, giving the shiny appearance of metals.



2.4 Recycling of Metals (self-reading) Metals are of prime importance to human needs. The most common and useful are Ni, Cr, Cu, Fe, Ag, Sn and Pb and their greatest use is in structural materials. With few exceptions, e.g. Mg from seawater and some Mn from the ocean floors, metals are produced commercially from subsurface ores. The common ores are oxides, sulfides, halides, silicates, carbonates and sulfates that are a finite resource. Hence recycling is a worthwhile process. For example, in the recycling of aluminium, used aluminium cans are shredded and pressed into blocks before being melted in a furnace. It requires 5% of energy to produce a tonne of recycled aluminium as it does to produce aluminium from bauxite ores. For every tonne of aluminium recycled, there is much less use of the raw materials and all the resources necessary to process them.

III. Ionic (Electrovalent) Bonding

(a) describe ionic (electrovalent) bonding, as in sodium chloride and magnesium oxide, including the use of „dot-and-cross‟ diagrams

(m) describe, in simple terms, the lattice structure of a crystalline solid which is: (i) ionic, as in sodium chloride, magnesium oxide

Ionic (Electrovalent) bonds are formed when one or more valence electrons are . from atoms of one element to another, resulting in the formation of positive and negative ions.

Ionic bond is the electrostatic attraction between two oppositely charged ions.

Ionic bonds are mostly formed between reactive metals (Gp I & II) and non-metals (Gp VI & VII), i.e. atoms with large electronegativity1 differences.

Group IV elements tend not to participate in ionic bonding since a lot of energy is required to

lose or gain 4 electrons in order to attain stable octet configuration.

3.1 ‘Dot-and-Cross’ Diagram (i) NaCl

Na +

1s22s22p63s1

Cl

1s22s22p63s23p5

[ Na ] [ Cl ]

1s22s22p6 1s22s22p63s23p6

(ii)

MgCl2

Mg +

1s22s22p63s2

2Cl

1s22s22p63s23p5

[ Mg ]2 2[ Cl ]

1s22s22p6 1s22s22p63s23p6

[Note: Strictly 'dot-and-cross' showing all valence electrons in atoms or ions, NO “circles” or

“shells” to be represented.]

1 Electronegativity is a measure of the relative attraction of an atom for electrons in covalent bonds (refer

to page 11).



3.2 Structure of Ionic Compounds The ions in an ionic compound attract ions of the opposite charge and repel ions of the same charge in every direction round them. This results in ionic compounds having a .where oppositely charged ions are arranged in a regular three-dimensional lattice structure. Hence, it has a giant ionic lattice structure.

The ions are arranged to allow the greatest number of contacts between oppositely charged ions without pushing ions of the same charge together.

3.3 Nature of the ionic bond Ionic bond is non-directional as an ion attracts an oppositely charged ion in all directions

without a preferred orientation. E.g. Na+ ion is attracted to all six Cl ions surrounding it in its

solid ionic lattice. The strength of ionic bonds is measured by the magnitude of lattice energy. Lattice energy is the energy when one mole of an ionic solid is formed from its constituent gaseous ions under standard conditions (298 K and 1 atm).

E.g. Li(g) + Cl(g) LiCl(s) H = -1030 kJ mol1

Na(g) + Cl(g) NaCl(s) H = -788 kJ mol1

K(g) + Cl(g) KCl(s) H = -701 kJ mol1

Mg2(g) + 2 Cl(g) MgCl2(s) H = -2326 kJ mol1

where

q+ & q charge of cation and anion respectively

r+ & r radius of cation and anion respectively

q q

r r

Fig. 4 Sodium chloride lattice structure In NaCl crystal, each Na+ ion is

surrounded by six Cl ions and each Cl

ion is surrounded by six Na+ ions. The

Na+ and Cl ions are held strongly by

electrostatic forces of attraction (electrovalent or ionic bonds).

Strength of ionic bonds Magnitude of Lattice energy (| lattice energy |) ||

rr

qq



Factors affecting Lattice Energy: (a) Charge of ion:

The more the ions, the stronger the attraction between the oppositely charged ions. This results in greater magnitude of lattice energy, stronger ionic bond.

(b) Size of ion:

The the ions, the smaller the interionic distance r+ + r, the closer they can approach each other in the solid lattice. This results in greater magnitude of lattice energy, stronger ionic bond.

[Note: Charge is a more important factor than radius.]

Question: Explain which of the following ionic compounds has stronger ionic bond. (a) sodium fluoride, NaF or sodium chloride, NaCl

Strength of ionic bonds | lattice energy | ||

rr

qq

The charge of all ions in NaF and NaCl are the , but the ionic radius of F

is than that of Cl. Thus, the magnitude of lattice energy of NaF is than

that NaCl, i.e. NaF has a stronger ionic bond than NaCl.

| lattice energy | of NaF and NaCl are 923 kJ mol1 and 788 kJ mol1 respectively. i.e. more energy is absorbed to weaken all ionic interactions in 1 mole of NaF to form its gaseous ions compared to overcoming the ionic interactions in 1 mole of NaCl to form its

gaseous ions.

(b) magnesium fluoride, MgF2 or sodium fluoride, NaF

Strength of ionic bonds | lattice energy | ||

rr

qq

Both compounds have the same anion, but the charge of Mg2+ is than that of Na+ and the ionic radius of Mg2+ is than that of Na+. Thus, |lattice energy| of MgF2 is than that of NaF, i.e. MgF2 has a stronger ionic bond than NaF.



3.4 Properties of Ionic Compounds Readily in polar solvents (e.g. H2O) but in non-polar solvents

(e.g. tetrachloromethane CCl4 and benzene C6H6)

Fig. 5 For an ionic compound to dissolve, the crystal lattice must be and this requires a great amount of energy. This amount of energy is offset by the release of energy caused by solvation (or hydration if the solvent is water). During solvation, polar solvents surround the ions by forming interactions. This bond formation releases sufficient energy to overcome the lattice energy and break down the ionic lattice structure, thus the solid dissolves. Ionic compounds do not dissolve in non-polar solvents since the interactions is too weak to overcome the strong electrostatic attraction between the ions in the crystal lattice. [Note: Not all ionic compounds dissolve in water, e.g. Al2O3 and AgCl are generally

insoluble. (Find out why in Chemical Energetics!)] http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/thermochem/solutionSalt.html

Good conductors of electricity in molten and aqueous state (dissolved in water)

In molten and aqueous states, the crystal lattice breaks down into ‘free ions’. These ions are . and act as charge carriers, and thus conduct electricity when a potential difference is applied. They are non-conductors when in solid state since ions are held tightly together in their lattice structure and not mobile.

High melting and boiling points (low volatility)

Due to strong electrostatic forces of attraction between oppositely charged ions, large amount of is required to separate the ions (weaken the ionic bond to break the lattice structure).

http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/thermochem/solutionSalt.html

http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/thermochem/solutionSalt.html



Hard and brittle

The ions are held by strong electrostatic attractions in a three-dimensional giant ionic lattice structure. When a force is applied, it causes slight displacement of a layer of ions. As a result, ions of like charge are brought together and the subsequent repulsion forces the crystal apart. The cleavage is possible because of the regular arrangement of ions in the crystal lattice.

IV. Covalent Bonding

(d) describe covalent bonding in terms of orbital overlap, giving σ and π bonds (g) explain the terms bond energy, bond length and bond polarity and use them to compare

the reactivities of covalent bonds

Covalent bond is the electrostatic force of attraction between the positive nuclei of two atoms and the negative bonding electrons shared between them.

A covalent bond is formed from the of a pair of electrons between atoms with high ionisation energies and similar electronegativities (i.e. between non-metallic atoms) to achieve a stable noble gas configuration. An electrically neutral group of atoms joined together by covalent bonds is called a molecule. Polyatomic ions contain a group of atoms joined together by covalent bonds and are electrically not neutral. There are many examples of covalent bonding: Elements - H2, O2, Cl2 Compounds - CO2, CH4, CH3CH3

Polyatomic ions - CH3+, CO3

2, NO2+, NO2

, NO3, SO3

2,SO42

4.1 Valence Bond Theory: Overlapping of Orbitals Covalent bonds are formed by the of orbitals. There are two types of bonds arising from different ways of overlapping of orbitals: a) Sigma Bond, σ

This bond is formed when the valence orbitals overlap (collinearly).

Fig. 6 An ionic lattice shatters when deformed.



The σ bond can be formed between head-on overlap of:

(i) s and s orbital e.g. in H2

+

(ii) s and p orbital e.g. in HCl

+

(iii) p and p orbital e.g. in Cl2

+

b) Pi bond, This type of bond arises from the overlap of p orbitals, producing a region of high electron density above and below the plane. Example: side-way overlap between p orbitals in O2

+

Though a bond has two regions of high electron density above and below the plane, it

constitutes only .

-bond is weaker than σ-bond because the degree of overlap of orbitals is lesser.

The first covalent bond formed is ALWAYS σ-bond. -bond can only be formed after σ-bond is formed. i.e. in any multiple bond, there is only one σ -bond and the rest are

-bonds.

4.2 Strength of covalent bonds The strength of a covalent bond is determined by the bond energy, which is the energy required to break one mole of covalent bond between two atoms in the gaseous state. Examples:

H2(g) → 2H(g) H = +436 kJ mol1

HCl(g) → H(g) + Cl(g) H = +431 kJ mol1

H measures strength of a covalent bond. The greater the bond energy, the stronger the bond.



Factors affecting strength of covalent bond: (a) Bond order – number of covalent bonds that exist between two atoms. The the bond order, the stronger the bond.

Covalent bond C-C C=C C≡C

Bond order 1 2 3

Bond energy / kJ mol-1 350 610 840

(b) Bond length – inter-nuclear distance between two covalently bonded atoms. Bond length

depends on the effectiveness of orbital overlap. The the orbital overlap, the shorter bond length and hence the stronger the covalent bond.

Bond Bond length / nm Bond energy / kJ mol-1

Cl – Cl 0.198 [ = 2 r(Cl) ] 244

Br – Br 0.228 [ = 2 r(Br) ] 193

(c) Bond polarity – the term 'polar' describes a covalent bond in which one atom had a partial

positive charge (+) and the other a partial negative charge (-). This arises due to difference in between the two atoms.

Electronegativity refers to the electron-pulling power of an atom for the bonded pair of electrons in a covalent bond. It is measured on a numerical scale with a value of 4.0 assigned to the most electronegative fluorine atom. Generally, it increases across a period and decreases down a group.

Fig. 7

A covalent bond is said to be polar or have ionic character if bonding atoms have different electronegativity values. The more polar the covalent bond, the stronger the ionic character in the covalent bond. This strengthens the bond as it gives additional interaction between two nuclei.

Bond Difference in E.N Bond energy/ kJ mol-1

N - H 0.9 390

O - H 1.4 460

F - H 1.9 562

In a non-polar bond between identical atoms, e.g N – N, Cl – Cl, F – F etc, there is equal

sharing of electrons between the atoms.



V. Shapes of Molecules and Polyatomic ions

(b) describe, including the use of „dot-and-cross‟ diagrams, (i) covalent bonding, as in hydrogen; oxygen; nitrogen; chlorine; hydrogen chloride;

carbon dioxide; methane; ethene (ii) co-ordinate (dative covalent) bonding, as in formation of the ammonium ion and in

the Al2Cl6 molecule. (c) explain the shapes of, and bond angles in, molecules such as BF3 (trigonal planar); CO2

(linear); CH4 (tetrahedral); NH3 (trigonal pyramidal); H2O (non-linear); SF6 (octahedral) by using the Valence Shell Electron Pair Repulsion theory

(e) predict the shapes of, and bond angles in, molecules analogous to those specified in (c)

Unlike metallic and ionic bonds (which are non-directional), covalent bonds are directional as shared electrons are localised (between 2 atoms). This gives rise to the shapes of covalent molecules and polyatomic ions. A pair of valence electrons involved in bonding is known as a of electrons. A pair of valence electrons not involved in bonding is called a of electrons. [Note: It is possible that lone / unpaired electrons may exist, i.e. not bonded and not in pairs, e.g. in NO2.]

5.1 Principles of Valence Shell Electron-Pair Repulsion (VSEPR) Theory Used in Predicting Molecular Shapes

The shape of a molecule is influenced by the number of lone pairs and bond pairs of electrons around the central atom. 1) Electron pairs would arrange themselves (in a three-dimensional space) as far apart as

possible to minimise mutual repulsion. (i.e. by keeping bond angles as large as possible) 2) Multiple bonds are treated as equivalent to single bonds as far as shapes of molecules

are concerned. 3) The extent of electron repulsion is:

lone pair - lone pair > lone pair - bond pair > bond pair - bond pair [Note: If present, lone or unpaired electron may be treated as a pair for the purpose of

bond angle and shape consideration.]



5.2 Summary of Shapes of Molecules and Polyatomic Ions

Lone

Pairs

No. of region

of electron

density

Electron

Pair

Geometry

Bonding

Pairs

Molecular

Geometry

Bond

angle

180o

120o

<120o

109.5o

107o

104.5o



5

6

90o;

120o

90o;

<120o

90o

180o

Bond

angle

Lone

Pairs

No. of region

of electron

density

Electron

Pair

Geometry

Bonding

Pairs

Molecular

Geometry

90o

90o

90o



5.3 Determining the Shape of Molecules ‘Dot-and-Cross’ Diagram

Step 1: Identify the central atom

The most electropositive (or least electronegative) atom is usually the central atom. H atom cannot be the central atom even if it is more electropositive. It is always an outer atom, because it can only form 1 bond.

The atom that is able to have an expanded octet (Period 3 onwards) or able to form many bonds.

Those that are not the central atom are the surrounding atoms. Step 2: Draw the outermost electrons in all atoms using only dots and crosses

(Different symbols used for neighbouring atoms)

The number of outermost electrons for each atom depends on the respective groups they are in. e.g.

C – Group IV 4 outermost electrons Cl – Group VII 7 outermost electrons

Step 3: Each surrounding atom should have at least 1 bond with the central atom

Ensure all surrounding atoms achieve octet configuration (8 electrons, except for H which is duplet) by forming bonds (can be single, double, triple or dative) with the central atom.

The central atom can have more or less or equal to 8 electrons.

Draw ‘.x’ to represent bond pair of electrons.

Draw ‘..’ or ‘xx’ to represent lone pair of electrons. Determine the shape of molecule and bond angles:

Step 4: Identify number of regions of electron density around central atom, which gives the electron pair geometry.

Step 5: Determine the arrangement of electron pairs (both bond and lone pair) around the

central atom. Step 6: Predict the molecular geometry or shape using the summary.

Molecules ‘Dot-and-cross’ diagram Structural formula Shape of molecule

O2

N2

HCl



Molecule ‘Dot-and-cross’ diagram Structural formula Shape of molecule

CO2

Bond angle:

BF3

Bond angle:

C2H4

Bond angle:

CH4

Bond angle:

NH3

Bond angle:

H2O

Bond angle:

PCl5

Bond angle:

SF6

Bond angle:



5.4 Co-ordinate or Dative Bonding

A co-ordinate (dative) bond is a covalent bond in which the shared pair of electrons is provided by only one of the bonded atoms.

One atom is the , the other is the . Once formed, a coordinate bond has the same characteristics as a covalent bond. The donor must have at least one lone pair of electrons in its outermost shell. ( A: ) The acceptor has must have at least one vacant orbital in its outermost shell. ( B ) The co-ordinate bond is represented as A: B Dative bonding plays an important role in the formation of molecular addition compounds and complex ions. 1. Molecular addition compound, BF3.NH3 molecule

Nx .

x.

H

H

x .H

B x.

x .

x .

F

F

F

. .....

... .

. .

. .....

+ Nxx

x .x.

H

H

x .H

B x.

x .

x .

F

F

F

. .....

... .

. .

. .....

..

2. Dimerisation of AlCl3 forming Al2Cl6

At temperature less than 200 oC, aluminium chloride exists as a , Al2Cl6, by forming dative bond between two AlCl3 monomers. On further heating, dimer dissociates to give two AlCl3 monomers.

http://www.yteach.co.uk/page.php/resources/view_all?id=electronegativity_polarity_variation_periodic_table_group_bond_t_page_21&from=search

Al

Cl Cl

Cl

Al

Cl

Cl

Cl





The following rules apply for the drawing of polyatomic ions. 1. For positive ion, electron is lost from the least electronegative atom. For negative ion, electron is gained by the most electronegative atom. 2. The charge is placed outside the bracket.

Substance ‘Dot-and-cross’ diagram Structural formula Shape of molecule

NH4

+

Bond angle:

H3O

+

Bond angle:

AlH4

Bond angle:



5.5 Exceptions to the Octet Rule Octet rule is sometimes not obeyed. (a) Molecules with an odd number of electrons (unpaired electron) e.g. NO, NO2, ClO2

N O N OCl

OOO

(b) Molecules with a central atom less than 8 valence electrons e.g. BeCl2, BF3, AlCl3

(c) Molecules with a central atom having more than 8 valence electrons (i.e. expanded

octet, applicable only for elements of period 3 and beyond) e.g. PCl5, SF6

Question: Why phosphorus can have more than 8 electrons (expanded octet) in its

valence shell e.g. in PCl5? Electronic configuration of P atom at ground state is 1s2 2s2 2p6 3s2 3p3

Electronic configuration of P atom in excited state is 1s2 2s2 2p6 3s1 3p3 3d1. From the valence electronic configuration of P, it has vacant low lying 3d orbitals, promotion / excitation of electrons from lower energy paired 3s orbital to higher energy 3d orbitals is possible. Since there are 5 unpaired electrons in excited state of P, it can form 5 single bonds with Cl in PCl 5.

Question: Can N (in Period 2 and in the same group as P) form NCl 5?

Electronic configuration of N atom at ground state is 1s2 2s2 2p3 N have vacant low lying d orbitals, hence electrons in paired 2s orbital cannot be promoted / excited. Since there are only 3 unpaired electrons in N atom, it can only form 3 single bonds with Cl. Hence NCl5 do not exist.

Be ClCl B

F

F FAl

Cl

Cl Cl

P

Cl

Cl

Cl

Cl

Cl

S

F

F

F

F

F

F



VI. Intermediate Bond Types

Fig. 8

6.1 Covalent Character of Ionic Bonds A bond is said to be purely ionic when there is complete transfer of electrons. However, in many cases, the positive charge on the cation in an ionic compound can attract electrons from an anion towards it. As a result, the electron clouds of the anions are distorted and this effect is known as . The ionic compound is thus said to exhibit a certain degree of covalent character.

Greater polarisation = greater degree of covalency.

Extent of polarisation depends on:

Polarising power of cation – The ability of the cation to attract electrons and distort an anion, which is favoured by high charge density.

where radius cationic

cation of charge density Charge

Polarisability of anion – The ease in which the electron cloud of an anion can be distorted, which is favoured by large size and high charge.

Question: Would NaF or NaI, show greater covalency in its bonding?

Both NaF and NaI, have the same Na cation while the anionic size of I is than F.

Therefore the electron cloud of I is than F. Hence, NaI shows greater

covalency in its bonding. Question: Would NaF or MgF2, show greater covalency in its bonding?

Both NaF and MgF2 have the same F anion. However, since the charge density of Mg2+

is (higher charge and smaller cationic size) than Na+, Mg2+ has a .than Na+

. Hence, MgF2 shows greater covalency in its bonding.

ionic bond

complete transfer of electrons

ionic bond with covalent character

partial transfer of electrons

covalent bond with ionic character

unequal sharing of electrons

covalent bond

equal sharing of electrons

increasing overlap of charge clouds

decreasing separation of charge



Question: How does covalency in ionic bond affect the nature of bonding in AlCl3? Although Al is a metal, its chloride is essentially covalent:

Al3+ is small in size and is highly charged. Thus, Al

3+ has a high charge density and

hence high polarising power.

Cl has a large electron cloud and is highly polarisable.

Hence, Al3+ polarises the electron cloud of Cl

to a large extent. This results in sharing of electrons and AlCl3 exists as a covalent compound (with ionic character).

6.2 Ionic Character of Covalent Bonds In a non-polar covalent bond, e.g. H2, each atom gets equal share of the bond pair of electrons. There is equal distribution of electron density. In a polar covalent bond, the more electronegative atom the bonding electrons to

itself (thus acquires a partial negative charge, -) and the other less electronegative atom will

acquire an equal partial positive charge, +. There is unequal sharing of electrons or uneven distribution of electron density. The more unequal the sharing, the more ionic character the covalent compound has. Greater difference in electronegativity = greater polarity in covalent bond.

VII. Dipole Moments and Polarity of Molecules

The + and - charge produce a dipole (a pair of separated charges of opposite signs), giving rise to a dipole moment. Dipole Moment: A vector quantity measuring the extent of polarisation (partial charge separation) in a covalent bond or molecule.

E.g. + - H─Cl

By convention, the dipole moment of a bond is denoted by an arrow (with a positive sign at one

end pointing from + to - end).

[Note: The + to - DO NOT imply existence of ions in the molecule. Ions carry full charges.]

7.1 Dipole Moment of a Molecule This is equal to the vector sum of the dipole moments of all the bonds present in the molecule.

A molecule is one where the overall dipole moment of the molecule is not zero. A molecule is one where the overall dipole moment of the molecule is zero (i.e. all dipole moments due to the polar bonds cancel each other) or there is no polar bond in the molecule.



7.2 Polarity of Molecules

To determine whether a molecule is polar or non-polar, consider the

1) bond polarity of all the bonds in the molecule and

2) shape of the molecule.

Examples i) H2O

O-H bonds are and H2O is in shape (or V-shaped).

Dipole moment do not cancel each other completely.

H2O is polar.

ii) CCl4

C-Cl bonds are and CCl4 is in shape.

Dipole moment cancel each another completely.

CCl 4 is non-polar.

[Note: Highly symmetrical molecules are usually non-polar.]

Question: Which of the following molecules are polar?

Br2 NF3 BCl3 CHCl3 C2H4



VIII. Intermolecular Forces of Attraction

(h) describe intermolecular forces (van der Waals‟ forces), based on permanent and induced dipoles, as in CHCl3(l); Br2(l) and the liquid noble gases

(f) describe hydrogen bonding, using ammonia and water as examples of molecules containing -NH and -OH groups

(m) describe, in simple terms, the lattice structure of a crystalline solid which is: (i) hydrogen-bonded, as in ice

(n) outline the importance of hydrogen bonding to the physical properties of substances, including ice and water

These are attractive forces between molecules or atoms and are important in explaining physical properties such as solubility, melting and boiling point. They are electrostatic in nature, and are very much weaker than the covalent bonds between atoms within the molecule. Three Main Types of Intermolecular Forces of Attraction

van der Waals’ Forces

Instantaneous dipole – induced dipole attraction

Permanent dipole – permanent dipole attraction

8.1 van der Waals’ forces (instantaneous dipole – induced dipole) attraction

Origin of instantaneous dipole – induced dipole attraction:

(a) At a particular instant, the of electrons may not be even or symmetrical for a non-polar molecule because electrons are in a state of continual motion.

(b) The resulting instantaneous dipole a dipole in the neighbouring molecules. (c) The electrostatic attraction between the temporary dipole and induced dipole is

known as van der Waals’ forces (instantaneous dipole – induced dipole) attraction. (d) Such attractions also exist for monoatomic elements such as the noble gases.

Fig. 9

Instantaneous dipole – induced dipole attraction exists between non-polar (and even polar) molecules i.e. ALL molecules.

Instantaneous dipole – induced dipole attraction are expected to be short-ranged and weakest among the three types of intermolecular forces for compounds with similar Mr.

http://www.yteach.co.uk/page.php/resources/view_all?id=intermolecular_force_matter_dipol_polar_non_polar_dispersion_hydrogen_bonding_t_page_14&from=search

Hydrogen bond





8.3 Factors affecting van der Waals’ forces of attraction

(a) Relative Molecular Mass

The higher the Mr, the electrons (larger the electron cloud) to be polarised and the the van der Waals’ forces (instantaneous dipole – induced dipole) attraction, therefore the higher melting point and boiling point.

Melting point/ oC Boiling point/ oC

F2 220 188

Cl 2 101 34

Br2 7 59 I2 114 184

(b) Shapes of molecules (applicable only for isomers)

Isomers have same molecular formula but different spatial arrangements of atoms. Example:

Molecular formula

Compound Name B.p. / oC

C5H12 Mr = 72

CH3CH2CH2CH2CH3

Pentane 36.3

C5H12 Mr = 72

CH3CH2CHCH3 | CH3

2-methylbutane 27.9

C5H12 Mr = 72

CH3 | CH3 C CH3

| CH3

2,2-dimethylpropane

9.5

Explanation:

- Boiling point of isomers is affected by the shape of molecules or degree of branching.

- 2,2-dimethylpropane is the most in shape (or has highest degree of branching) while pentane is a straight chain molecule.

- Hence the surface area of contact between molecules of 2,2-dimethylpropane is the while highest in pentane.

- This results in the van der Waals’ forces (instantaneous dipole - induced dipole) attraction in 2,2-dimethylpropane and in pentane.

- energy required to weaken the vdW forces in 2,2-dimethylpropane while greatest amount of energy is required to overcome the vdW forces in pentane.

- Hence boiling point in 2,2-dimethylpropane and highest boiling point in pentane.

- 2-methylbutane is less branched than 2,2-dimethylpropane but more branched than pentane. Thus its boiling point is between that of 2,2-dimethylpropane and pentane.



8.2 van der Waals’ forces (permanent dipole – permanent dipole) attraction

Origin of permanent dipole – permanent dipole attraction:

(a) Polar molecule has a dipole moment (non-zero dipole moment).

(b) The electrostatic attraction between the permanent dipoles is known as van der Waals’ forces (permanent dipole – permanent dipole) attraction.

Fig. 10 Permanent dipole – permanent dipole attraction

Permanent dipole – permanent dipole attraction exists between molecules.

As the magnitude of dipole moment increases, strength of permanent dipole – permanent dipole attraction increases.

[Note: Generally, permanent dipole – permanent dipole attraction is more significant than instantaneous dipole – induced dipole attraction between polar molecules.] Example (to see that id-id attraction can be more significant than pd-pd attraction)

Mr Shape Polarity Boiling point (oC)

SO2 64.1 Bent Polar -10

SO3 80.1 Trigonal planar Non-polar 45

SO2 is a molecule and the intermolecular forces of attraction between SO2 molecules is van der Waals’ forces (permanent dipole – permanent dipole) attraction. SO3 is a molecule and the intermolecular forces of attraction between SO3 molecules is van der Waals’ forces (instantaneous dipole – induced dipole) attraction. Generally, van der Waals’ forces (permanent dipole – permanent dipole) attraction is stronger than van der Waals’ forces (instantaneous dipole – induced dipole) attraction. However, SO3 has a higher boiling point than that of SO2. WHY??? SO3 has a Mr as compared to SO2, hence there are more electrons to be polarised in SO3. Therefore, the van der Waals’ forces (instantaneous dipole – induced dipole) attraction between SO3 molecules become than the van der Waals’ forces (permanent dipole – permanent dipole) attraction between SO2 molecules, accounting for the higher boiling point of SO3. [Note: For molecules with similar size of electron cloud (Mr): strength of hydrogen bonds > van der Waals‟ forces (permanent dipole – permanent dipole) attraction > van der Waals‟ (instantaneous dipole – induced dipole) attraction]



Question: How do we know which reason to quote???

To analyse the physical data given and quote the APPROPRIATE factor to explain it. This is a MUST under Assessment Objective (k).

8.4 Hydrogen bonding

A hydrogen bond is the electrostatic attraction between:

o a hydrogen atom covalently bonded to a highly electronegative atom (limited to F, O & N)

and

o a lone pair of electrons of another highly electronegative atom (limited to F, O & N) that comes from another molecule.

The H atom has a + charge and the electronegative atom has a - charge (a permanent dipole).

http://programs.northlandcollege.edu/biology/Biology1111/animations/hydrogenbonds.swf

8.4.1 Factors affecting hydrogen bond

compound formula b.p. / oC

ammonia NH3 -33

water H2O 100

hydrogen fluoride HF 19.5

(a) Extent of hydrogen bonding (number of hydrogen bond per molecule)

H2O has a higher boiling point than NH3. The hydrogen bond between the H2O molecules is more extensive than that

between NH3 molecules. This is because H2O, on average, has two hydrogen bonds per molecule whereas

NH3 has one hydrogen bond per molecule.

Hydrogen bond

http://programs.northlandcollege.edu/biology/Biology1111/animations/hydrogenbonds.swf



(b) Strength of hydrogen bond (dipole moment of H-X bond)

HF has a higher than boiling point than NH3. The hydrogen bond between HF molecules is stronger than the hydrogen bond

between NH3 molecules. This is because H – F bond is more polar than N – H bond.

[Note: The extensiveness of hydrogen bond between molecules predominates.]

8.4.1 Intermolecular and intramolecular hydrogen bonding

Compound

2-nitrophenol

4-nitrophenol

Structure

O

H

N

O

O

intramolecular hydrogen bonding

O

H

NOO

O

H

NO

O

intermolecular hydrogen bond

Boiling points / C 216 259

Type of intermolecular forces to be overcome during boiling

The -OH and -NO2 groups within the molecule are able to form intramolecular hydrogen bonds (due to the close proximity).

This the extent of intermolecular hydrogen bonds to be formed between the molecules (less extensive hydrogen bonding between the molecules).

The -OH and -NO2 groups within the molecule are NOT able to form intramolecular hydrogen bonds

Thus, only intermolecular hydrogen bonds between molecules

extent of intermolecular hydrogen bonds to be formed between the molecules is not affected (more extensive hydrogen bonding between the molecules).

Relative amount of energy required to overcome the intermolecular hydrogen bond

lesser

higher

[Note: Covalent bonds are NOT broken during boiling for simple molecular compounds.]

- +

+

-



8.4.2 Influence of hydrogen bonding on properties of substances (a) Dimerisation of carboxylic acids in non-polar organic solvent e.g. CCl 4

The Mr of ethanoic acid in CCl4 is found to be 120 instead of 60 due to dimerisation. The

acid is polar whereas the organic solvent is non-polar. The solute - solvent interactions are than hydrogen bonds formed among the acid molecules. The acid molecules therefore dimerise via hydrogen bonding.

hydrogen bond

CH3 C

O

O H

O

C

O

CH3

H

In aqueous solvents, ethanoic acid forms hydrogen bond with water. No dimerisation occurs. [Note: Solubility requires intermolecular forces of solute and solvent to be compatible.]

(b) Unusual properties of water

In ice, the water molecules are held in fixed position by hydrogen bonds. The presence of two hydrogen atoms and two lone pairs of electrons in each water molecule enable 4 hydrogen bonds to be formed per water molecule (or on average, two hydrogen bonds per water molecule). These give rise to an open, crystalline structure (resulting in large volume) with hexagonal holes as shown below. Hence, ice is less dense than water.

Structure of water in the liquid and solid state:

-

- +

+

Water

Ice



(c) Anomalous trends in boiling points of Group V, VI, VII hydrides

Question: Explain the trend of boiling points of these hydrides down the group.

(i) Consider boiling points of Group IV hydrides: CH4, SiH4, GeH4 and SnH4

All are simple covalent molecules with weak van der Waals’ forces (instantaneous dipole - induced dipole) attraction among molecules. As the number of electrons to be polarised from CH4 to SnH4, the intermolecular forces of attraction increases and the energy required to overcome them increases. Thus, the boiling point increases from CH4 to SnH4.

(ii) Consider boiling points of H2O, NH3 and HF with hydrides of the respective

Groups.

H2O, NH3 and HF are simple covalent molecules with intermolecular hydrogen bonding. However, the corresponding hydrides are only able to form weak van der Waals' forces between the molecules. Since hydrogen bonds are than van der Waals' forces, more energy is required to overcome the hydrogen bonds between the molecules. Thus, H2O, NH3 and HF have exceptionally high boiling points.

IX. Properties of Molecular Substances

(m) describe, in simple terms, the lattice structure of a crystalline solid which is: (ii) simple molecular, as in iodine (iii) giant molecular, as in graphite; diamond

9.1 Simple Covalent Molecules (a) Low melting and boiling point

During melting and boiling, energy is required but there is no increase in temperature. Energy is used to overcome forces of attraction between molecules, i.e. intermolecular forces of attraction. NO strong covalent bond is broken when melting and boiling occur in simple molecular compounds.



Strong covalent bond

Weak Intermolecular forces

i.e hydrogen bond

Strong

covalent bond

H2O(g)

H2O(l)

Less energy is needed to weaken the weak van der Waals’ forces of attraction or hydrogen bonds, therefore these substances have lower melting and boiling points.

(b) Solubility in polar and non-polar solvents

Polar molecules are soluble in solvents. Molecules which can form hydrogen

bonds with water are generally soluble in water (compatible solute solvent interactions). Non-polar molecules are soluble in solvents, as non-polar solvent molecules

are able to penetrate and break down the solute structures.

Solubility requires intermolecular forces of attraction between solute and solvent to be compatible. The solute will dissolve in the solvent when energy released on forming intermolecular forces of attraction between solute and solvent is sufficient to compensate for the energy required to break the solvent – solvent interactions and solute – solute interactions.

(c) Non-conductor of electricity in solid or molten state

The absence of delocalised mobile electrons and mobile ions as charge carriers indicate the inability of simple molecular substances to conduct electricity in solid and molten state.

If the molecule is polar, e.g. HCl, it can dissolve in water and ionise (forming H3O

+ and

Cl), thus able to conduct electricity in aqueous state.

9.2 Giant Covalent Molecules Giant covalent molecules are made up of atoms (same or different elements) held by strong covalent bonds. (a) Large amount of energy is required to break the covalent bonds between atoms

(b) Insoluble in most solvents

(c) Non-conductors of electricity: Electrons in these molecules are in the covalent bonds, are not mobile to act as charge carriers and hence they are not able to conduct electricity (exception: graphite).



9.2.1 Examples of Giant Covalent Molecules Graphite

The structure consists of flat, parallel layers of arranged carbon atoms. Each C atom is covalently bonded to three other C atoms in the same layer. The fourth

electron in each C atom forms electron cloud with neighbouring C atoms by the overlapping of p orbitals.

Conducts electricity in directions to planes as delocalised electrons (due to electron cloud) move along layers (but not between layers).

Layers are held by van der Waals’ forces of attraction. Thus, the layers slide over each other easily making graphite soft for use as pencil lead and high-temperature lubricant.

High melting point as strong covalent bonds within layers have to be broken.

Diamond

Structure consists of a 3-dimensional network of covalently bonded C atoms arranged

tetrahedrally in space. Each C atom is bonded to four other C atoms. Non-conductor of electricity since all electrons are , no charge carrier present. High melting point as strong covalent bonds have to be broken.

Silicon(IV) Oxide, SiO2

Each silicon atom is covalently bonded to four oxygen atoms tetrahedrally and each oxygen is bonded to two silicon atoms.

Non-conductor of electricity since all electrons are bonded. High melting point as strong covalent bonds have to be broken. [Note: The diagram is a simplification of the actual structure of silicon dioxide. In reality, the

"bridge" from one silicon atom to its neighbour isn't in a straight line, but via a "V" shape due to presence of 2 lone pairs on each oxygen atom.]



Comparison between Diamond and Graphite

S/No. Diamond Graphite

1. C atoms form strong covalent bonds with each other in 3-D.

C atoms forms covalent bonds within a layer but the intermolecular forces of attraction between layers is van der Waals’ forces due to

instantaneous dipole induced dipole attraction.

2. High melting point (strong covalent bonds are directed in space)

High melting point (strong covalent bonds within the layers)

3. High density, 3.5 g cm-3 (short bond length)

Lower density, 2.3 g cm-3 (distance between layers relatively larger)

4.

Low electrical conductivity: all the four valence electrons are localised and are not mobile to act as charge carriers throughout the lattice.

High electrical conductivity. One of the four valence electrons of each C atom is delocalised: mobile electrons act as charge carriers throughout each layer.

5. Rigid and hard structure Soft and greasy feeling (used as lubricant and pencil lead)

Remaining assessment objectives encompassing the whole chapter and not attributed to any particular sub-topic: (j) describe, interpret and/or predict the effect of different types of bonding (ionic bonding;

covalent bonding; hydrogen bonding; other intermolecular interactions; metallic bonding) on the physical properties of substances

(k) deduce the type of bonding present from given information (l) show understanding of chemical reactions in terms of energy transfers associated with the

breaking and making of chemical bonds (under chemical energetics) (o) suggest from quoted physical data the type of structure and bonding present in a

substance.

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