2014 Chem Bonding (FINAL)(1)

Nanyang Junior College H2/H1 Chemistry JC1 2014

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Chemical Bonding

Lecturer: Ms Grace Leong and Ms Joanne Low _______________________________________________________________________________

Content 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 Metallic bonding Bonding and physical properties The solid state

Learning Outcomes 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 and pi 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


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(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


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1. Introduction to Chemical Bonding

Only the noble gases exist as free atoms at room temperature and pressure. Atoms of all other elements are held together by chemical bonds.

Chemical bonds are electrostatic forces of attraction that bind atoms together in an element or compound. The bond forming process is exothermic. Energy is given out such that the overall energy of the system is lowered and stability is attained.

Only valence electrons (electrons in the outermost principal quantum shell) participate in bonding. These electrons maybe completely transferred from one atom to another or shared between two or more atoms. This means that in most cases, atoms tend to gain, lose or share electrons until their outermost principal quantum shell is full and a noble gas electronic configuration is achieved.

The bonding in an element or compound will affect its structure. The structure of the element or compound will in turn determine its physical properties, for example, melting / boiling points, electrical conductivity, solubility and so on.

Bonding Structure Physical properties


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Overview

References 1. Chemistry for Advanced Level. Peter Cann 2. Chemistry The Molecular Nature of Matter and Change. Silberberg

CHEMICAL BONDS

Structure

Giant Metallic Lattice

Giant Ionic Lattice

E.g. hydrogen, water

Structure

E.g. diamond, graphite, silicon dioxide

Giant Molecular Structure

Structure Structure

Explains physical properties: - m.p / b.p - electrical conductivity - solubility, etc

Metallic bond Electrostatic

attraction between cations and sea of

delocalised electrons

Bonding

Covalent bond Electrostatic attraction between nucleus and

shared pair of electrons

Bonding

Ionic bond Electrostatic

attraction between cations and

anions

Bonding

1. van der Waals 2. pd-pd 3. H-bonding

Intermolecular forces

METALLIC E.g. iron, copper, sodium, calcium

COVALENT E.g. hydrogen, water, silicon dioxide, diamond

IONIC E.g. sodium chloride, magnesium oxide

Particles

cations surrounded by a sea of delocalised valence electrons

cations and anions Particles

molecules Particles

atoms

Particles

Simple Molecular Structure


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2. Metallic Bonding and Structure

2.1 Introduction

Structure Metals exist as a giant metallic lattice

Particles of cations surrounded by a sea of delocalised electrons,

Bonding held together by strong metallic bonds.

Definition: Metallic bonds are strong electrostatic attraction between the cations and the sea of delocalised electrons in the giant metallic lattice.

Only valence electrons are delocalised. (Delocalised: electrons are free to move throughout the whole structure and are not bound to any one particular atom.) The sea of delocalised electrons prevent the cations from repulsion.

2.2 Factors affecting metallic bond strength

Number of delocalised valence electrons per cation

The larger the number of delocalised valence electrons per cation, the stronger the electrostatic attraction between the more highly-charged cations and sea of delocalised electrons, the stronger the metallic bond.

Eg: Na < Mg < Al

Size of cation

The smaller the cation, the stronger the electrostatic attraction between the cations and sea of delocalised electrons, the stronger the metallic bond.

Eg: Na > K > Rb

Sea of

delocalised

electrons

Cations


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2.3 Physical properties of metals

2.3.1 Melting and boiling points

(Always link mp/bp to energy required to overcome...)

Element Melting point / C

Sodium, Na 98

Magnesium, Mg 651 Aluminium, Al 660

Potassium, K 63

Rubidium, Rb 39

Metals have generally high melting and boiling points. A lot of energy is required to overcome the strong metallic bond between the cations and sea of delocalised electrons in the giant metallic lattice during melting and boiling.

The stronger the metallic bond, the higher the melting and boiling point.

Can you account for the melting points of Na vs Mg vs Al? Na vs K vs Rb?

2.3.2 Electrical conductivity

Electrical conductivity requires the presence of mobile charge carriers. Mobile charge carriers can be moving ions or electrons. (Always link electrical conductivity to mobile charge carriers.)

Metals are good conductors of electricity due to the sea of delocalised electrons in the metallic lattice. The sea of delocalised electrons act as charge carriers to conduct an electric current around the electric circuit. The cations are held in fixed positions and do not act as mobile charge carriers.

Electrons move from the negative terminal to the positive terminal of the electrical circuit. Note: By convention, current flow is in the opposite direction to electron flow!

+


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2.3.3 Heat conductivity (independent reading)

When heat is applied to one end of a metal, the kinetic energy of the electrons and cations increase. Kinetic energy is transmitted through the delocalised electrons and hence rapidly heats up the cooler end. Energy can also be passed via particle vibrations of the cations.

2.3.4 Surface luster and reflectivity (independent reading)

When light falls on a metal surface, electrons absorb light energy and are excited into unfilled orbitals. The electrons return to lower energy states by emitting the absorbed energy as light. This gives metals a shiny, mirror-like surface.

2.3.5 Malleability and ductility

Malleable: Able to be hammered into sheets (opposite: brittle) Ductile: Able to be pulled into wires

A metallic bond is strong but flexible. When an external force is applied to a metal, the cation layers slide past one another and the sea of delocalised electrons move to prevent repulsion between the cations. In other words, the cations aquire new neighbours but are still held together by the sea of delocalised electrons.

If metals are malleable and ductile, why are silver and gold jewellery, or metallic pots and pans so hard?

Force applied

along this plane

Cation layers slide

past each other


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2.4 Uses of metals (independent reading)

Metals are useful materials because they are malleable and ductile, are good conductors of electricity and have high melting points. However, metals are a finite resource.

Researchers studying supplies of copper, zinc and other metals have determined that the earths supply of metals is unable to meet future demand. Furthermore, the process of mining and refining metals is costly.

It is therefore more environmentally friendly and cost effective to recycle metals. For instance, recycling aluminium requires less than 1% of the energy than that of the whole process required to obtain aluminium from its ore.

Uses of aluminium and copper Properties Uses

Alu

min

ium

Light and strong, especially when alloyed Aircraft parts, boats, cars

Corrosion-resistant When exposed to air, a layer of Al2O3 forms on the surface, preventing Al from further oxidation by O2 in the air.

Window frames

Good conductor of heat, non-toxic Cooking ware Light, non-toxic, resistant to corrosion Cans and foils for packaging food

Copp

er Good conductor of electricity Electrical wires

Good conductor of heat Cooking ware Resistant to corrosion Water pipes

QR code: Read about a Yale study on the depletion of metals.


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2.5 Alloys (independent reading)

An alloy is a metal with the addition of one or more other elements. Examples of alloys are steel (iron and carbon) and brass (copper and zinc).

The introduction of a 2nd element disrupts the orderly arrangement of the metal cations. The metal cations are no longer able to slide past one another easily. This strengthens the material so that it can be used for a wider range of applications.

Alloys are mixtures, not compounds. The 2nd element is not chemically bonded to the metal, and is also not added in a fixed proportion.

Uses of Common Alloys

Alloy Parent Element Alloying Element Characteristics Uses

Bronze Copper Tin Hard and strong Ship propellers

Brass Copper Zinc Corrosion resistant

Castings; hardware items; boilers in ships

Steel Iron Carbon Hard and strong Bridges, buildings

Metal ions

Second

element


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3. Ionic Bonding

3.1 Introduction

Structure Ionic compounds exist as a giant ionic lattice

Particles of cations and anions

Bonding held together by strong ionic bonds.

Definition: Ionic bonds are strong electrostatic attraction between the cations and anions in the giant ionic lattice.

Ionic bonds are usually formed between a metal and non-metal (also see Section 5). The atom that loses electrons (usually a metal) becomes a positive ion or cation, while the atom that gains electrons (usually a non-metal) becomes a negative ion or anion. The cations and anions formed usually have the electronic configuration of a noble gas, ie. octet (ns2np6).

Example: Dot and cross diagrams for ionic compounds MgO

AlF3

Na2O

Ca3N2

Giant ionic lattice of sodium chloride


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3.2 Factors affecting ionic bond strength

The strength of ionic bonds is related to the lattice energy of the ionic compound.

Definition of lattice energy: The energy released when 1 mole of a pure ionic solid is formed from its constituent gaseous ions under standard conditions. (Will be covered in the upcoming chapter on Chemical Energetics.)

Eg: Na+ (g) + Cl- (g) NaCl (s)

The magnitude of the lattice energy depends on the product of the ionic charge and sum of the ionic radii: Lattice energy

| q+ q- |r

+ + r-

Charge of ions

The greater the charge of the ions, the stronger the electrostatic attraction between oppositely charged ions, the greater the lattice energy and the stronger the ionic bonds.

Eg: NaCl vs MgO

Size of ions

The bigger the ions, the weaker the electrostatic attraction between oppositely charged ions, the smaller the lattice energy and the weaker the ionic bonds.

Eg: NaCl vs KBr

3.3 Physical properties of ionic compounds

3.3.1 Melting and boiling points

All ionic compounds are solids at room temperature as they have very high melting and boiling points.

A lot of energy is required to overcome the strong ionic bond between the cations and anions in the giant ionic lattice during melting and boiling.

The stronger the ionic bond, the higher the melting and boiling point.

Compound Melting point / C Sodium chloride 801

Magnesium oxide 2852 Potassium bromide 734 Copper(II) sulfate 600


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Can you account for the melting points of NaCl vs MgO? NaCl vs KBr? MgO vs CuSO4?

Polyatomic ions (also see Section 4.4) A polyatomic ion contains 2 or more atoms bonded together covalently and has an overall charge (there are additional or missing electrons). Common examples include the sulfate (SO42-), nitrate (NO3-), carbonate (CO32-) and ammomium (NH4+) ions.

3.3.2 Hard and brittle

Ionic compounds are hard as every ion is strongly attracted to the oppositely charged ions around it.

They are also brittle. When a little stress is applied, the ions get displaced and the ions with similar charge come together, repelling each other. The ionic lattice falls apart.

3.3.3 Solubility (also see Section 8)

Ionic compounds are usually soluble in water.

Water molecules will interact with the cations and anions through ion-dipole interactions, releasing energy. If the amount of energy released through these interactions is greater than the energy required to overcome the ionic bonds between the cations and anions in the giant ionic lattice, the ionic compound will be soluble in water.

-

-

-

-

+

+

+

+

+

+

+ +

-

-

-

-

+ +

+

+

+

+

+

+

Stress

applied along

this plane

Like charges

repel, solid

falls apart

ionic lattice

solvated cation

solvated anion


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However, there are ionic compounds that are insoluble in water, eg AgCl. In this case, we say that the amount of energy released during the formation of ion-dipole interactions is not enough to overcome the ionic bonds between the cations and anions in the giant ionic lattice.

Ionic compounds are usually insoluble in organic solvents. Organic solvents are carbon-containing compounds such as hexane, propanol and dichloromethane.

The energy released from the weak interactions formed between the solvent molecules and ions are not enough to overcome the ionic bonds between the cations and anions in the giant ionic lattice.

3.3.4 Electrical conductivity

Ionic compounds are non-conductors in the solid state but are good conductors in the molten or aqueous state.

Recall: Mobile charge carriers are required for electrical conductivity.

In the solid state, the ions are held in the giant ionic lattice by strong electrostatic attraction and can only vibrate about a fixed position. They are not mobile.

In the molten state or aqueous state, the lattice is broken down, the ions become mobile and are able to act as mobile charge carriers to conduct electricity.

Aqueous sodium chloride

Bulb lights up

Solid sodium chloride

Bulb does not light up

Molten sodium chloride

Bulb lights up

heat


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4. Covalent Bonding

4.1 Introduction

There are 2 types of covalent molecules the simple covalent molecule and the giant covalent molecule. A molecule with a few atoms covalently bonded together is a simple covalent molecule (has a simple molecular structure), while a molecule with tens of thousands of atoms covalently bonded together is a giant covalent molecule (has a giant molecular structure).

Definition: Covalent bonds are strong electrostatic attraction between the nuclei of atoms and their shared pair of electrons.

A covalent bond is formed when 2 atoms, A and B, approach each other such that the nucleus of atom A attracts the electron of atom B and the nucleus of atom B attracts the electron of atom A.

Recall that electrons in atoms are found in atomic orbitals. When the atomic orbitals overlap during the formation of the covalent bond, a molecular orbital is formed.

Imagine 2 H atoms approaching each other.

At 1, the atoms are infinitely far apart and do not interact with each other. The potential energy of the system is approximately 0.

As we go from 1 2, the distance between the atoms decrease. The electrons of A are attracted to the nucleus of B and vice versa and the energy is lowered.

At 2, the energy is at a minimum. The distance between the nuclei at 2 is optimum and is the covalent bond length. If the 2 atoms go even closer (from 2 3), the potential energy of the system rises again, because of repulsion between the 2 positively charged nuclei.

The bond strength/energy is the energy required to break the bond and separate the atoms back to infinity. (Also see Section 4.3.)

orbital overlap

2x 1s orbitals molecular orbital


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Electronegativity (also see Section 5)

Definition: The ability of an atom in a covalent bond to attract the bonding electrons.

When atoms of the same element form a covalent bond, the bonding electrons are equally shared between the atoms. The electron cloud is symmetrically distributed and this is a non-polar covalent bond.

H H Cl Cl C H

Image source: http://vigorouschemist.hubpages.com/hub/Chemical-Bonding-and-Electronegativity

(*For convenience, the electronegativity of C / H / P / B can be taken to be approximately equal.)

When atoms of 2 different elements form a covalent bond, the bonding electrons are not equally shared between the atoms. This is because each atom has a different electronegativity. Generally, the electronegativity of an atom increases across a period and decreases down the group. The electronegativity is related to the effective nuclear charge of the atom.

The more electronegative atom will attract the bonding electrons more strongly. This will result in the more electronegative atom acquiring a partial negative charge and the other atom acquiring a partial positive charge.

This separation of charge results in a dipole (represented by ) and the formation of a polar covalent bond.

H Cl H F I Cl

Always show the partial charges (+/-) when drawing a dipole!


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4.2 Sigma and pi bonding

Depending on how the atomic orbitals overlap, there can be 2 different types of covalent bonds formed.

4.2.1 Sigma bond

A sigma () bond is formed when atomic orbitals overlap head-on.

The shared pair of electrons occupy the space between the nuclei.

There can only be 1 sigma bond between 2 atoms (as there is no other way for another head-on overlap to take place). In other words, all single bonds are sigma bonds. Also note that s-orbitals can only overlap head-on.

Case 1: Head-on overlap of two s-orbitals

e.g. 1s orbital of H + 1s orbital of H H H bond

s-orbital s-orbital bond formed

Case 2: Head-on overlap of s-orbital and p-orbital

e.g. 1s orbital of H and 2p orbital of F H F bond

s-orbital p-orbital bond formed

Case 3: Head-on overlap of two p-orbitals

e.g. 3p orbital of Cl and 3p orbital of Cl Cl Cl bond

p-orbital p-orbital bond formed


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4.2.2 Pi bond

A pi (pi) bond is formed when atomic orbitals overlap side on.

The shared pair of electrons occupy the space above and below the nuclear axis. There is no electron density found along the nuclear axis.

Case 4: Side-on overlap of two p-orbitals

e.g. 2p orbital of O and 2p orbital of O in O2

p-orbital p-orbital pi bond formed

A pi bond can only form when the 2 atoms are already held together by a sigma bond. Pi bonds are weaker than sigma bonds due to less efficient overlap. All multiple bonds contain pi bonds.

Bond Order Name Number of shared electron pairs

Type of bonds formed

1 Single bond 1 1

2 Double bond 2 1 , 1 pi

3 Triple bond 3

Example: Bonding in O2 and N2

We know that there is a double bond between O atoms in O2 (O = O).

The double bond consists of 1 bond and 1 pi bond. It is formed based on the overlap of atomic orbitals.

Electronic configuration of

O: 1s2 2s2 2px2 2py1 2pz1

There is a (2p 2p) overlap and a pi(2p 2p) overlap.

Nuclear axis


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Electronic configuration of N: 1s2 2s2 2px1 2py1 2pz1

There is a ________ bond between N atoms in N2, consisting of ___________ and _________ bonds.

There is a (2p 2p) overlap and two pi(2p 2p) overlaps.

4.3 Factors affecting bond energies

Energy is taken in during bond breaking and released during bond formation.

Definition: Bond energy is the energy required to break 1 mole of gaseous covalent bonds under standard conditions (also see Chemical Energetics).

The larger the bond energy, the stronger the covalent bond, the less easily it breaks during chemical reactions.

In order of importance, the bond energy depends on the Bond order (generally, a triple bond > double bond > single bond) Bond length (generally, shorter bonds are stronger) Bond polarity (polar bonds are stronger than expected)

Bond energies can be found in the Data Booklet.

4.3.1 Bond order

The higher the bond order between any 2 atoms, the greater the degree of orbital overlap, the stronger the bond, the higher the bond energy.

E.g. C C (840 kJ mol-1) vs C = C (610 kJ mol-1) vs C C (350 kJ mol-1)

Example: The bond energy of C C is 350 kJ mol-1 but the bond energy of C = C is 610 kJ mol-1. Why is the bond energy of the double bond less than twice that of the single bond?

A single bond consists of 1 ____ bond. A double bond consists of 1 ____ bond and 1 ____ bond.

A ____ bond is weaker than a ____ bond due to less effective overlap.


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4.3.2 Bond length

Definition: Bond length is the internuclear distance between the 2 bonding atoms.

Bond length is dependent on the size of the atoms that are bonded.

1 pm = 1 x 10-12

m

The shorter the bond length (when bond order is the same), the greater the degree of orbital overlap, the stronger the bond, the higher the bond energy.

E.g. Cl Cl (244 kJ mol-1) vs Br Br (193 kJ mol-1) vs I I (151 kJ mol-1)

Example: Why is the bond energy of the F F bond (158 kJ mol-1) lower than expected?

The F F bond length is very short. The F atoms are pulled very close together and this leads to repulsion between the lone pairs on the F atom, weakening the F F bond.

4.3.3 Bond polarity

Polar bonds are stronger than expected. The presence of partial charges in the bond results in extra attraction.

Example:

Bond Bond energy / kJ mol1 H H 436 I I 151

Br Br 193 H I 299

H Br 366

The bond energy of H I is close to the average of H H and I I bonds (294 kJ mol-1) while the BE of H Br is larger than the average of the BE of H H and Br Br (315 kJ mol-1).


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The difference in electronegativity for H and I is not as large as compared to H and Br.

H I is relatively non-polar. Hence, the BE is close to the average of H H and I I which are non-polar.

H Br is polar. Hence, the BE is larger than the average of H H and Br Br due to the extra attraction between the opposite partial charges on H and Br.

4.4 Dot-and-Cross Diagrams

Dot-and-cross diagrams are used to represent the valence electrons of individual atoms in the covalent molecule or polyatomic ion.

Alternate the usage of dots and crosses between adjacent atoms. No other symbols besides dots and crosses can be used. There is no need for a key to explain the source of the electrons. All valence electrons must be shown, not just the bonding electrons.

General guide for drawing dot-and-cross diagrams:

1. Identify the central atom based on the number of bonds it can form. (E.g. H can only form 1 bond and cannot be the central atom.) The element with fewer atoms is usually the central atom. Arrange the other atoms (terminal atoms) around the central atom.

2. For polyatomic ions, draw a square bracket around the dot-and-cross diagram and write the overall charge on the top right hand corner outside the bracket.

If the ion is negatively charged, add the corresponding number of extra electrons to the more electronegative atom. The distribution of the extra charge should be spread out if possible, i.e. if there is 2 charge, distribute the 2 extra electrons to 2 different atoms.

If the ion is positively charged, remove the corresponding number of electrons from the less electronegative atom.

3. Work from the terminal atoms, ensuring that the terminal atoms achieve noble gas configuration by forming the necessary number of bonds with the central atom. Electron pairs between atoms are known as bond pairs.

4. Add the remaining electrons (total no. of valence e minus e involved in bonding with terminal atoms) as lone pairs to the central atom. Unshared electron pairs are known as lone pairs.

5. Check if the central atom has achieved octet (not always necessary, many possible exceptions, see Sections 4.4.1 to 4.4.3) o If the central atom is from Period 3 or above, it can have more than 8 electrons around it. o If the central atom is from Period 2, it cannot have more than 8 electrons around it. If it

does, consider changing a double bond to dative bond.


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Example: Dot and cross diagrams of covalent molecules / polyatomic ions Water

Carbon dioxide

Ammonia

Methane, CH4

NO2+

CO32-

4.4.1 Exception 1: Electron deficient species

Electron deficient molecules have fewer than 8 electrons around the central atom. Gaseous compounds containing either aluminium, beryllium or boron as the central atom are often electron deficient.

BeCl2

BH3


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4.4.1 Exception 2: Expansion of octet

Many molecules and ions have more than 8 valence electrons around the central atom. The central atom is able to accommodate the additional electrons because it has empty and energetically accessible 3d orbitals to accommodate the extra electrons. Generally, only non-metals from Period 3 onwards (Si) can expand their octet.

Vacant orbitals in quantum shells higher than the valence shell cannot be used for expansion of octet since the energy difference is too high (not energetically accessible).

Phosphorous pentachloride, PCl5

Sulfur dioxide, SO2

Why do you think SF6 exists but not SCl6?

Expansion of octet is also limited by the atoms size. 6 of the smaller F atoms can fit around the S atom, but not the larger Cl atoms due to steric hindrance.

4.4.3 Exception 3: Odd electron species (radicals)

Some molecules have an odd number of valence electrons so they cannot possibly have all the electrons paired. These molecules are known as radicals and are very reactive. Usually, the central atom carries the odd number of valence electrons.

Nitrogen monoxide, NO

Nitrogen dioxide, NO2

Radicals can dimerise (2 molecules combine) by forming a covalent bond via their unpaired electrons under certain conditions.


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2NO N2O2

2NO2 N2O4

4.5 Dative bonding

For some covalent bonds, both the electrons in the bonding pair are provided by only 1 of the bonded atoms. This type of bonding is known as dative bonding.

A dative bond is no different from any other covalent bond. It can form within a molecule/ion or between 2 particles. For the latter, one species must have a lone pair available for donation and the other must have be electron deficient and have empty, energetically accessible orbitals to accept the pair of electrons.

Carbon monoxide, CO

Ammonium ion, NH4+

NH3.BF3

Dimerisation of AlCl3 via dative bonding


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4.6 Lewis structures and structural formulae

Besides using dot-cross diagrams, Lewis structures can also be used to represent the bonding within molecules.

In a Lewis structure, each bond pair is represented with a line, and lone pairs are represented with 2 dots. Dative bonds are represented by a pointing from the donor atom to the acceptor atom.

E.g. water carbon dioxide carbon monoxide

Molecules can also be represented as their structural formulae, with electron pairs between atoms represented with lines. There is no need to show the lone pairs when drawing structural formulae.

E.g. carbon dioxide


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4.7 Shapes of molecules

Simple covalent molecules and polyatomic ions have definite shapes based on the arrangement of atoms. The shapes and structures of the molecules are important because they determine the physical and chemical properties of the substances.

Molecular shapes can be predicted by the Valence Shell Electron Pair Repulsion (VSEPR) theory. Note that in VSEPR, double bonds, triple bonds or dative bonds are treated as one electron pair.

4.7.1 VSEPR theory

VSEPR theory states that:

1. Electron pairs (bond pairs (bp) and lone pairs (lp)) arrange themselves in space as far as possible from each other to minimise mutual repulsion.

2. Lp lp repulsion > lp bp repulsion > bp bp repulsion.

(A bond pair is attracted by 2 nuclei and is situated approximately between 2 nuclei. A lone pair, however, is attracted by only 1 nucleus. It is hence nearer to the nucleus and exerts stronger repulsion on electron pairs, resulting in distorted bond angles.)

3. Repulsion between electron pairs depends on the electronegativities of the atoms concerned.

Based on the first two rules, all molecules with the same number of bond pairs and lone pairs should have exactly the same bond angles. However, this is not true, as the repulsion between electron pairs is also dependent on the relative electronegativity of the central and bonding atoms. The closer the electron density is to the central atom, the greater the repulsion, and the larger the bond angle.

NH3 (107) PH3 (94)

NH3 and PH3 have the same number of bond pairs and lone pairs and hence, shape.

However, N is more electronegative than P and pulls the bonding electrons closer to itself (the central atom). There is greater repulsion between the bonded electron pairs, leading to a larger bond angle in NH3.


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Total no. of electron pairs

No. of

bond pairs

No. of

lone pairs

Example Structural formula

or dot cross diagram

Shape, structure, bond angle

2

2 0

BeCl2

Linear, 180

3

3 0

BF3

Trigonal planar, 120

2 1

SO2

Bent, 118

4

4 0

CH4

Tetrahedral, 109

3 1

NH3

Trigonal pyramidal, 107

2 2

H2O

Bent, 105


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No. of

bond pairs

No. of

lone pairs




5

5 0

PCl5

Trigonal bipyramidal, 90 (Axial equatorial) and 120 (equatorial

equatorial)

4 1

SF4

See-saw

3 2

ClF3

T-shape

2 3

XeF2

Linear


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No. of

bond pairs

No. of

lone pairs




6

6 0

SF6

Octahedral, 90

5 1

BrF5

Square pyramidal

4 2

XeF4

Square planar


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How to determine the shapes of molecules or polyatomic ions using VSEPR theory:

1. Draw the appropriate dot-cross diagram / lewis structure / structural formula.

2. Determine the electron distribution based on the total number of electron pairs about the central atom arrange them as far apart from each other as possible to minimize repulsion.

3. Lone pairs contribute to the electron distribution (and hence, bond angle) about the central atom, but only atom arrangements contribute to the final shape of the molecule. Subtract off a vertex for every lone pair present to determine the final shape.

Example: Determine the shape of Cl2O

2 bond pairs, 2 lone pairs Shape: Bent

SO32-

_____ bond pairs, _____ lone pair Shape: Trigonal pyramidal


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4.7.2 Cases with more than 1 central atom

We apply VSEPR theory to each central atom in the molecule for molecules with more than 1 central atom.

Example: Determine the shape about each central atom in H2O2

2 bond pairs, 2 lone pairs

Shape: bent about each O

C2H4

_____ bond pairs, _____ lone pairs

Shape: __________________ about each C

C2H6


Shape: __________________ about each C

N2O4


Shape: __________________ about each N


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5. Intermediate Bonding

We have assumed that compounds are either completely ionic (electrons transferred) or completely covalent (electrons shared). In actual fact, pure ionic or pure covalent bonds seldom exist. Instead, we can view ionic vs covalent bonding as a spectrum, with most bonds taking on intermediate characters.

The type of bond formed depends on the difference in electronegativities (see Page 15) of the elements involved in bonding.

QR code:

Watch an animation to see the difference between ionic bonds,

non polar covalent bonds and polar covalent bonds.

Pure ionic bond:

Electrons are completely

transferred with no

electron density

between the cation and

anion

Pure covalent bond:

Electrons are shared

equally due to no

difference in

electronegativity

between the 2 atoms

Polar covalent bond:

Electrons are shared

unequally due to a slight

difference in

electronegativity

between the 2 atoms

Ionic bond with covalent

character:

Distortion (polarisation)

of anions electron cloud

by cation of with high

polarising power

Bonding Spectrum Large difference in electronegativity

No difference in

electronegativity


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5.1 Ionic bonds with covalent character

In a purely ionic compound, both the cation and the anion have perfectly spherical electron clouds and there is no electron density between them.

However, if the cation has a high polarising power, it might attract the anions electron cloud so strongly that it causes distortion. There is some degree of sharing of electrons between the 2 ions.

The greater the extent of distortion, the more significant the covalent character. In the extreme case, the sharing of electron clouds can be so significant that it becomes a predominantely covalent bond, e.g. AlCl3 and BeCl2.

The factors affecting distortion are

Polarising power of cation

The higher the charge density (ionicchargeionicsize

) of the cation, the higher its polarising power.

Polarisibility of anion

The larger the size of the anion, the further the electron cloud is from the attraction of its nucleus, the higher its polarisibility.

Generally, bonding tends to be ionic if the size of cation is large (low polarising power) and anion is small (low polarisibility), or the charges on both cation and anion are small.

5.2 Polar covalent bonds (See Page 15)

When atoms of the same electronegativity (e.g. H H, Cl Cl) are covalently bonded together, the bonding electrons are shared equally and the electron cloud is symmetrical. This is a non-polar bond.

When atoms of different electronegativity are covalently bonded together (e.g. H Cl, C Br), the bonding electrons are shared unequally, with the electron density skewed towards the more electronegative atom. This sets up a partial charge distribution (+ / -), known as a dipole. There is slight ionic character in the polar bond.


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5.3 Polar molecules

A polar molecule is one that has a net dipole moment. The overall polarity of a molecule depends on 2 criteria.

1. The presence of polar bonds

If there are no polar bonds in the molecule, the molecule is definitely non-polar.

2. The arrangement of the polar bonds

If there are polar bonds in the molecule, check whether the dipoles cancel each other out. If they cancel out and there is no net dipole, the molecule is non-polar.

Molecules with more than 1 equivalent polar bond that are symmetrically arranged about the central atom are non-polar (e.g. SF6, PCl5, CF4). The dipoles cancel out and the net dipole moment is 0.

A molecules polarity will affect its physical properties.

Example: State the shape of the following molecules and determine if they are polar. Methane, CH4

4 bp, 0 lp, tetrahedral Non polar (no polar bonds)

Chloromethane, CH3Cl

4 bp, 0 lp, tetrahedral

Carbon dioxide

___ bp, ___ lp, ______________

SF2

___ bp, ___ lp, ______________


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6. Simple molecular structures

In this section, we focus on covalent molecules with simple molecular structures (e.g. O2, HCl, CO2, H2O) and their physical properties.

Structure (Name of molecule) has a simple molecular structure

Bonding with (select the appropriate intermolecular force)

Particles between molecules

6.1 Introduction to intermolecular forces

Intermolecular forces are the forces between simple covalent molecules. They are weak forces (as compared to covalent bonds, ionic bonds and metallic bonds). Intermolecular forces affect the physical properties (e.g. melting and boiling point) of the substance.

There are strong covalent bonds between the atoms. The covalent bonds affect the chemical properties of the substance.

Note that no covalent bonds are broken during the melting and boiling of substances with simple molecular structures! E.g. when solid iodine sublimes to form gaseous iodine, the I I bond does not break. When water boils, it still exists as water molecules, not H and O atoms.

Solid iodine Liquid iodine Gaseous iodine


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There are 3 main types of weak intermolecular forces between simple molecules: Van der waals Permanent dipole permanent dipole Hydrogen bonds

6.2 Van der Waals forces

Van der waals (vdW) forces are present between all molecules (and atoms for the noble gases). They are the most significant intermolecular forces between non-polar molecules.

What are van der Waals forces? In an atom or molecule, the electrons are constantly moving. At any moment, there can be an asymmetrical electron density around a molecule, resulting in an instantaneous dipole. This instantaneous dipole will induce another dipole on its neighbouring molecule.

The electrostatic attraction between these fluctuating dipoles are known as van der Waals forces, and enable molecules to be attracted to one another.

(You might see the term instantaneous dipole induced dipole forces in some textbooks. It is okay to use either term.)

These attractive forces are very weak as the dipoles are continuously turned on and off as the electron density fluctuates.

QR code:

Read more about Johannes Diderik van der Waals,

the scientist who first postulated the presence of an intermolecular force.


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6.2.1 Factors affecting the strength of van der Waals forces

1. Number of electrons per molecule (can be inferred from Mr)

The greater the number of electrons per molecule, the more polarisable the electron cloud, the stronger the van der Waals forces.

Example: Explain why the boiling point of Group VII elements increases down the group.

Element No of electrons per molecule Boiling point / C Appearance at room temperature

Chlorine 34 -35 Yellowish-green gas

Bromine 70 59 Reddish-brown liquid

Iodine 106 184 Black solid

Chlorine, bromine and iodine all have simple molecular structures with weak van der Waals forces between molecules. (standard intro for all bonding questions involving physical properties)

The vdW forces increase from chlorine to bromine to iodine as the number of electrons increase leading to a more polarisable electron cloud.

More energy is required to overcome the stronger vdW forces, leading to a higher boiling point.

2. Shape of molecule (only for molecules with similar number of electrons per molecule)

The shape of the molecule may also influence the strength of the van der Waals forces. This is usually applied to structural isomers in Organic Chemistry. Elongated molecules have larger surface areas of contact than spherical molecules.

The larger the surface area of the molecule, (the greater the area of contact), the stronger the van der Waals forces.

-


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Example: Explain the variation in boiling points for the following isomers of C5H12. Isomer pentane methylbutane 2,2-dimethylpropane

Mr 72 72 72

Structure CH3CH2CH2CH2CH3

b.p. / 0C 36 28 9

All the isomers have simple molecular structure with weak van der Waals forces between molecules.

Pentane has the most straight-chain structure and 2,2-dimethylpropane has the most spherical shape due to its highly branched structure.

The vdW forces increase from 2,2-dimethylpropane to methylbutane to pentane as the surface area of contact between molecules increases.

More energy is required to overcome the stronger vdW forces, leading to a higher boiling point.

6.3 Permanent dipole permanent dipole interactions

Polar molecules (e.g. HCl, ICl) have a permanent net dipole. The oppositely charged ends of the dipoles attract each other. This attraction is known as permanent dipole permanent dipole (pd-pd) interactions. Covalent bonds are 100 times stronger than permanent dipole permanent dipole interactions.

There are also van der Waals forces between polar molecules but for most cases, we can consider the predominant intermolecular force to be pd-pd interactions.

+ - + - + -

+ - + - + -

+ - + - + -


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Example: Explain the variation in boiling points between the substances. Substance Mr Net Dipole moment, / D Boiling point / 0C Ethane, CH3CH2CH3 30 Negligible -89 Hydrogen sulfide, H2S 34 0.97 -61 Methanal, HCHO 30 2.7 -19 Bromine, Br2 160 0 59

For molecules with similar number of electrons, permanent dipole permanent dipole interactions are stronger than van der Waals forces.

E.g. ethane vs methanal

Ethane and methanal have simple molecular structure. Ethane is a non-polar molecule with weak van der Waals forces between molecules. Methanal is a polar molecule with permanent dipole permanent dipole interactions between molecules.

More energy is required to overcome the stronger pd-pd interactions than vdW forces, so methanal has a higher boiling point than ethane.

For molecules with similar number of electrons, the larger the permanent dipole moment, the stronger the permanent dipole permanent dipole interactions between molecules.

E.g. hydrogen sulfide vs methanal

Hydrogen sulfide and methanal have simple molecular structure with permanent dipole permanent dipole interactions between molecules.

The net dipole is larger in methanal than hydrogen sulfide as O is more electronegative than S. More energy is required to overcome the stronger pd-pd interactions between methanal so it has a higher boiling point.

We cannot compare pd-pd vs vdW for molecules with very different number of electrons (but questions are not likely to be set in this context. Usually, data will be given and you will be asked to suggest a reason for the trend.)

E.g. hydrogen sulfide vs bromine

(standard intro) + the vdW forces between bromine molecules is more significant than pd-pd interactions between hydrogen sulfide molecules + (energy req)


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6.4 Hydrogen bonds

A hydrogen bond is the electrostatic attraction between a hydrogen atom bonded to a small nitrogen, oxygen or fluorine atom and a lone pair on the nitrogen, oxygen or fluorine atom of a neighbouring molecule.

Therefore, the conditions for hydrogen bonding between molecules are:

H bonded to F, O or N on molecule Lone pair on F, O or N of neighbouring molecule

A hydrogen bond is an extreme case of a permanent dipole permanent dipole interaction. It is 10 50 times stronger than van der Waals forces and 5 10 times stronger than pd-pd interactions (but still 10 20 times weaker than a covalent bond).

A hydrogen bond is strong because the high electronegativities of F, O or N leads to relatively large dipole moments. Furthermore, the hydrogen atom is small and bare, and able to get close to the lone pair of electrons for strong attraction.

Examples of hydrogen bonding between molecules: (Always show the lone pair, partial charges and hash lines when drawing hydrogen bonds.)

Between water molecules

Between hydrogen fluoride molecules

Between ammonia molecules

H F H F


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Show the hydrogen bonding interaction between water and oxygen gas:

Example: This graph shows the trend in boiling points of the Group IV, V, VI and VII hydrides.

Why does the boiling point increase from CH4 to SnH4 (and as we go down each group)?

The Group IV hydrides have _________________________ with ___________________________

between _______________.

As we go down the group, the ________________________________ increases, and the

________________________ is ___________________________.

More _______________ is needed to overcome the stronger vdW forces, leading to a higher boiling

point.


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Why are the boiling points of NH3, H2O and HF considerably higher than that of CH4 (and higher than expected), even though the relative molecular masses are similar?

CH4, NH3, H2O and HF have __________________________. There are

_______________________ between CH4 molecules but _____________________ between NH3,

H2O and HF molecules.

More _______________ is required to overcome the stronger _______________ than

_________________________, so NH3, H2O and HF have considerably higher boiling points.

Why is the boiling point of H2O > HF > NH3?

All 3 have simple molecular structure with hydrogen bonding between molecules.

There is an average of 2 hydrogen bonds per molecule in H2O but only an average of 1 hydrogen bond per molecule in HF and NH3.**

More energy is required to overcome the more extensive H-bonding in H2O as compared to HF and NH3, hence the highest boiling point.

F is more electronegative than N so the dipole formed is larger.

More energy is required to overcome the stronger hydrogen bonds between HF molecules so HF has a higher boiling point than NH3.

**The average number of hydrogen bonds per molecule is the total number of sets of hydrogen atoms + lone pairs in the molecule. H atoms Lone pairs Sets NH3 3 1 1

H2O HF


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6.4.1 Intermolecular vs Intramolecular hydrogen bonding

An intermolecular hydrogen bond is formed during the interaction of 2 or more molecules.

An intramolecular hydrogen bond between a H bonded to a F/O/N and another F/O/N atom present in the same molecule. It usually occurs in organic compounds with suitable functional groups in close proximity.

If intramolecular hydrogen bonding is present, intermolecular hydrogen bonding will be less extensive (because the lone pairs and H atoms are already used up during intramolecular hydrogen bonding).

Example: Account for the difference in boiling points of 2-nitrophenol and 4-nitrophenol.

Compound Boiling point / C

214

290

2-nitrophenol and 4-nitrophenol both have simple molecular structures with hydrogen bonding between molecules.


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There is intramolecular hydrogen bonding in 2-nitrophenol as the OH and NO2 groups are close to each other (and hence, less intermolecular hydrogen bonding), but not in 4-nitrophenol.

More energy is required to overcome the more extensive hydrogen bonding between 4-nitrophenol

6.4.2 Dimerisation through hydrogen bonding

When the relative molecular mass of organic acids such as ethanoic acid is measured in the gas phase or in a non-polar solvent, the relative molecular mass appears to be about twice the theoretical values, suggesting the presence of dimers.

Mr in aqueous solution Mr in gaseous phase Mr in benzene

Ethanoic acid, CH3COOH 60 120 120

2 acid molecules can dimerise through hydrogen bonding.

(Recall: besides hydrogen bonding, how else can molecules dimerise? Why do they dimerise that way?)

Dimerisation does not take place in aqueous solution as the acid molecules can form hydrogen bonds with surrounding water molecules instead. Note that the solvent is usually taken to be in large excess.

Show the hydrogen bonding interaction between water and ethanoic acid:


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6.4.3 Structure and density of ice

Liquids are usually less dense than solids as the particles in a liquid are packed less closely together. However, this is not the case for H2O as ice floats on water! This can be explained by hydrogen bonding.

In the liquid state, the molecules have enough energy to overcome the hydrogen bonds and roll and slide over one another. There is no regular arrangement of molecules.

When water freezes, the water molecules no longer have enough energy to roll and slide over one another. Hydrogen bonds form permanently and the water molecules move away from each other as they freeze into an open lattice structure to maximize the effects of hydrogen bonding.

In the open lattice structure, each oxygen atom is surrounded by 4 hydrogen atoms in a tetrahedral arrangement. 2 of these hydrogen atoms are covalently bonded and the other 2 hydrogen atoms are attracted via hydrogen bonding.

The volume of water increases, resulting in a decrease in density. This phenomenon allows aquatic life to survive under ice.

1st

QR code:

Animation about waters polarity,

how it forms ion-dipole interactions with a solute such as NaCl 2

nd QR code: Animation of water molecules freezing into ice.

3rd

QR code: 3D structure of the open lattice of ice


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7. Giant molecular structures

In this section, we focus on substances with giant molecular structures giant covalent molecules consisting of thousands of atoms held together by strong covalent bonds throughout the structure to give a 3D lattice. Giant covalent molecules are also known as macromolecules.

Structure (Name of molecule) has a giant molecular structure

Particles with atoms

Bonding held together by strong covalent bonds

Common examples of substances with giant molecular structures include graphite, diamond, silicon, silicon dioxide (SiO2), boron nitride (BN) etc.

Diamond vs graphite Diamond and graphite are allotropes. Allotropes are 2 or more forms of the same element, in which the atoms are arranged in different ways. Other allotropes of carbon include the carbon fullerenes and nanotubes etc.

Graphite Diamond

Stru

ctu

re

Shap

e ab

ou

t C at

om

Each carbon atom is bonded to 3 other carbon atoms in the same layer.

Trigonal planar (120) about each carbon atom

Each carbon atom is bonded to 4 other carbon atoms.

Tetrahedral (109) about each carbon atom.


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Graphite Diamond De

scrip

tion

There is a network of layers of hexagonal rings.

There is one free 2p electron per carbon atom that is delocalised along the layer (not between layers).

There are strong covalent bonds within the layer, and weak Van der Waals forces of attraction between layers.

There is a network of carbon atoms with tetrahedral geometry.

All valence electrons are used in bonding.

There are strong covalent bonds throughout the molecule.

Pro

pert

ies Soft and black

High melting and boiling point (why?) Electrical conductor along the layers (why?)

Hard and transparent High melting and boiling point (why?) Non-electrical conductor (why?)

Uses

Pencil lead Dry lubricant (why?) Inert electrodes

Gemstones Tips of cutting, grinding and polishing tools, drill bits (why?)

To think about: Why do substances with giant molecular structures have high melting and boiling points? Why is graphite a good conductor of electricity but not diamond? Why is graphite used as a dry lubricant? Why is diamond used in drill bits?

Silicon dioxide (SiO2, also known as silica) has a similar structure to diamond. Each Si atom is covalently bonded to 4 other O atoms in a tetrahedral arrangement and each O atom is covalently bonded to 2 other Si atoms in a bent arrangement.

What kind of physical properties do you think silicon dioxide has?

QR code 1:

Read more about the fullerenes.

The discovery of the fullerenes won 3 scientists the Nobel Prize in Chemistry in 1996.

QR code 2:

An interview with Nobel Chemist Harry Kyoto about his discovery.


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8. Predicting Solubility

Generally, a major factor in determining whether a solution can be formed is the relative strength of the interactions between the solute and solvent particles.

The energy released during the interactions formed between the solute and solvent should be comparable to or more than the energy taken in to overcome the forces between the solute particles and between the solvent particles.

(Before mixing)

(After mixing)

To explain why something is soluble, simply state the types of interactions between the solute and solvent, e.g.

Ammonia is soluble in water as it can form hydrogen bonds with water

Sodium chloride is soluble in water as it can form ion-dipole interactions with water

Tetrachloromethane is soluble in benzene as it can form van der Waals interactions with benzene

A

A

B

B

A

B

overcome overcome

form


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To explain insolubility, discuss relative amounts of energy involved during the interactions between the solute and solvent vs interactions between the solute particles and between the solvent particles, e.g.:

Diamond is insoluble in water as the energy released during the formation of van der Waals forces with water is not enough to overcome the strong covalent bonds between carbon atoms (and hydrogen bonds between water molecules).

Oil is insoluble in water as the energy released during the formation of van der Waals forces with water is not enough to overcome the van der Waals forces between oil molecules and hydrogen bonds between water molecules.

Oil will float on water as it has a lower density than water. The mixture can be separated using a separating funnel.

Octanol is insoluble in water even though it can form hydrogen bonds with water. Its predominantly forms van der Waals forces with water due to its long, non-polar hydrocarbon chain. The energy released during formation of these van der Waals forces is not enough to overcome the hydrogen bonds between water molecules (and the van der Waals forces between octanol).

Long non-polar hydrocarbon chain

Tap


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8.1 Application: How soaps work (independent reading)

It is not possible to clean an oily surface (which attracts dirt) using just water, as oil and water are immiscible. However, the cleaning process is made a lot easier with the aid of soap and detergent.

Surfactants are organic molecules that have a water soluble component, and an oil soluble (water insoluble) component. Surfactants are often used in soaps and detergents as they can interact with both oil and water so that oil and dirt can be trapped, and removed and rinsed away by water.

The surfactant, sodium stearate, CH3(CH2)16COO- Na+ is a major component of many bar soaps.

The non-polar tails of the surfactant can interact with the oil and dirt via van der Waals forces to form a micelle (an aggregate of surfactant molecules dispersed in a liquid), trapping the oil and dirt inside the micelle.

The ionic heads of the surfactant forms ion-dipole interactions with the surrounding water molecules, and the micelle will be washed off by water, bringing the oil and dirt along with it.

QR code:

Animation about how soaps work.

Note that some phrases used in the video e.g. hydrophobic, hydrophilic etc

are biology terms and although correct, should not be used in Chemistry.

Stick to non-polar and polar for chemistry.

Long non-polar hydrocarbon chain (tail)

Ionic head

Common (simplified)

representation of

surfactant


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9. Summary

Complete this table on your own as a form of revision. You might also want to create your own mind map to summarise the topic.

Ionic Bonding Metallic Bonding Covalent Bonding Definition (electrostatic attraction between)

Representation

Diagram:

Dot-cross:

Diagram: Dot-cross:

Usually formed between

Examples

Exceptions

Factors affecting strength (list down simple keywords you must include in explanation for each factor)

Structure (structure, particles, forces of attraction)

To be completed in Page 51.

MP / BP

Electrical conductivity

Solubility


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COVALENT BONDING

Structure Simple molecular structure Giant Molecular Structure Particles

Forces of attraction

Van der Waals Permanent dipole permanent dipole

Hydrogen bonding

Usually formed between (indicate type of molecules)

List examples:

Examples

Factors affecting strength (list down simple keywords you must include in explanation for each factor)

MP/BP

Electrical conductivity

Solubility

Documents

2014 Chem Bonding (FINAL)(1)