unit 1: the core principles of chemistry - … · unit 1: the core principles of chemistry ... o...

Dr. Catherine Tan

1

+6018 – 375 2833 www.chemistryboutique.com

Unit 1: The Core Principles of Chemistry (Edexcel International Advanced Level Chemistry)

Energetic Enthalpy Changes, H

o Chemical reactions are accompanied by energy changes.

o Energy changes can be exothermic ( H < 0) and endothermic ( H >0)

o Enthalpy, H is an indication of a substance’s total energy content and it cannot be

measured directly but enthalpy change H is measurable

Exothermic Reactions

o It gives out heat to the surrounding.

o Therefore, surrounding temperature rises as the heat content of the system decrease.

o H is negative value.

E.g. C(s) + O2(g) CO2(g) H = -393.4kJ/mol

Endothermic Reaction

o It absorbs heat from the surrounding.

o Thus, the surrounding temperature decrease as the heat of the system rises.

o H is positive value.

Dr. Catherine Tan

2


E.g. H2(g) + I2(g) 2HI(g) H = +51.9kJ/mol

o Exothermic reactions are energetically more favorable than endothermic reactions because a

system with lower heat content is more stable.

o Hence, most chemical reactions are exothermic.

o The more negative a H value, the more stable is the system.

Enthalpy Change of Reaction

o It is defined as the heat change when the reaction takes place between the masses of the

reagents indicated by the stoichiometric equation for the reaction.

o Enthalpy change depends on:

– Temperature

– Physical states of the reactants and products

– Pressures of gaseous reactants and products

– Concentration of solution

o Enthalpy changes are stated under standard conditions:

– Pressure: 1atm

– Temperature: 25oC

– Substance in its most stable physical form

– Enthalpies of elements in their standard states are taken to be zero.

o e.g. N2(g) H = 0 kJ/mol

o A thermochemical equation gives the amount of reactants and products (measured in moles)

as well as the quantity of energy involved.

4NH3 + 3O2 2N2 + 6H2O H = - 1260kJ/mol

NH3 + 3/4O2 1/2N2 + 3/2H2O H = -1260/4 kJ/mol

2N2 + 6H2O 4NH3 + 3O2 H = +1260 kJ/mol

Enthalpy change of Formation, Hf

o It is defined as the enthalpy change when 1 mole of the compound is formed from its

elements under standard conditions (25oC, 1atm).

o Hf is usually negative but some are positive (e.g. Hf for oxides, NOx).

o Hf of elements in its standard state is zero.

o Hf is often used to predict the stability of a compound relative to its constituent elements.

o e.g. H2(g) + 1/2O2(g) H2O(l) Hf = -286kJ/mol

Dr. Catherine Tan

3


C(graphite) + O2(g) CO2(g) Hf = -393kJ/mol

C(graphite) + 2H2(g) CH4(g) Hf = -75kJ/mol

K(s) + Mn(s) + 2O2(g) KMnO4(s) Hf = -813kJ/mol

• Hf <0, then compound is energetically more stable than its constituent elements.

• Hf >0, then compound is energetically less stable than its constituent elements.

Using Enthalpies of Formation to Calculate Enthalpy of Reactions

o Consider combustion of propane

o C3H8 (g) + 5O2 (g) 3CO2 (g) + 4H2O (l)

o Calculate the enthalpy change of combustion using the information below.

o Given Hf C3H8 (g) = -104kJ mol-1, Hf CO2 (g) = -394kJ mol-1, Hf H2O (l) = -286kJ mol-1

Solution

H = Hf products - Hf reactants

= (3x-394) + (4x-286) - (-104)

= -2222kJ mol-1

Enthalpy change of combustion, Hc

o It is defined as the enthalpy change when 1 mole of a substance is completely burnt in

oxygen under standard condition (25oC 1atm).

o Hc is always negative, as heat is always evolved in the combustion.

o Hc can be used to give the energy values of fuels and foods.

Example:

S(s) + O2(g) SO2(g) Hc = -297kJ/mol

CH4(g) + 2O2(g) CO2(g) + 2H2O(l) Hc = -890kJ/mol

C2H4(g) + 3O2(g) 2CO2(g) + 2H2O(l) Hc = -1411kJ/mol

Dr. Catherine Tan

4


Using Enthalpies of Combustion to Calculate Enthalpy of Reactions

Q: Given

C (s) + O2 (g) CO2 (g) Hf/c = -394kJ mol-1

H2 (g) + 1/2O2 (g) H2O (l) Hf/c = -286kJ mol-1

C3H8 (g) + 5O2 (g) 3CO2 (g) + 4H2O (l) Hc = -2200kJ mol-1

Calculate the H for the following reaction:

• 3C (s) + 4H2 (g) C3H8 (g)

Solution

Method 1

Remain equation 1 x3

• 3C (s) + 3O2 (g) 3CO2 (g) H = (-394)x3 kJ mol-1

Remain equation 2 x4

• 4H2 (g) + 2O2 (g) 4H2O (l) H = (-286)x4 kJ mol-1

Reverse equation 3

• 3CO2 (g) + 4H2O (l) C3H8 (g) + 5O2 (g) H = 2200kJ mol-1

Add up three equations

• 3C (s) + 4H2 (g) C3H8 (g)

• H = (-394x3) + (-286x4) + (+2200)

= -126kJ mol-1

Dr. Catherine Tan

5


Solution

Method 2 (Using enthalpies combustion)

• 3C (s) + 4H2 (g) C3H8 (g)

• H = Hc reactants - Hc products

= (-394x3) + (-286x4) - (-2200)

= -126kJ mol-1

Enthalpy change of hydration, Hhyd

o It is defined as the enthalpy change when 1 mole of gaseous ions is dissolved in a large

amount of water under standard conditions (25oC, 1atm).

o e.g. Na+(g) + aq Na+(aq) Hhyd = -406kJ/mol

Cl-(g) + aq Cl-(aq) Hhyd = -366kJ/mol

o Thus, the enthalpy change of hydration of NaCl is the enthalpy change that accompanies the

hydration of 1 mole of both the gaseous ions Hhyd is always negative, as the heat is

produced when bonds are formed between the ions and the dipoles on the water molecules.

o The hydration energies of the ions depend on the charge and size of the ions.

o The higher the charge and the smaller the size, the greater is the hydration energy.

Dr. Catherine Tan

6


Enthalpy change of solution, Hsol

o It is defined as the enthalpy change when 1 mole of a substance dissolves in such a large

volume of solvent that addition of more solvent produces no further heat change under

standard conditions (25oC, 1atm).

o Hsol can be positive or negative.

o Hsol is very positive, compound is insoluble in water.

o Hsol is negative, compound is soluble in water

Dr. Catherine Tan

7


Enthalpy change of atomisation, Hat

o It is defined as the enthalpy change when an element or a compound is converted into 1 mole

of atoms under standard conditions (25oC, 1atm).

o Hat is always positive, because energy must be absorbed to pull the atoms far apart and to

break all the bonds between them.

o The enthalpy of atomisation is not the same as the enthalpy of vaporisation.

o When an element is vaporised, the gas particles are usually not separate atoms.

Enthalpy change of neutralisation, Hn

• It is defined as the enthalpy change when 1 mole of water is formed in the neutralisation

between an acid and an alkali, the reaction being carried out in aqueous solution under

standard conditions (25oC, 1atm).

• Hn is always negative.

• Calculation of Hn can be measured by mixing the solutions of acids and alkalis in a

calorimeter and measuring the rise in temperature.

Dr. Catherine Tan

8


Neutralisation between Strong Acid and Strong Base

• Enthalpy of neutralisation between strong acid and strong base is almost constant(-

57.3kJ/mol).

HX(aq) H+(aq) + X-(aq)

MOH(aq) M+(aq) + OH-(aq)

H+(aq) + OH-(aq) H2O(aq) Hn = -57.3kJ/mol

Neutralisation between weak acid and weak base

• Since weak acids and weak bases are slightly ionised in aqueous solution.

• Thus, the enthalpy change of neutralisation involving weak acid and base is always

smaller than -57.3kJ/mol.

Electron affinity

o It is the energy change associated with the formation of an anion from the gaseous atom.

o The first electron affinity of an element represents the energy released when an electron is

added to an atom in the gaseous state.

o The second electron affinity is the energy absorbed for the process.

Bond Energy

o It is the energy required to break one mole of a covalent bond between two atoms in the

gaseous state.

o Bond breaking is endothermic whereas bond forming is exothermic.

o Since chemical reaction involving bond breaking followed by bond forming, hence the

enthalpy change of reaction is the energy difference between bond breaking and bond

forming.

Dr. Catherine Tan

9


o e.g. Combustion of alkanes is always exothermic. This is due to the energy released on

making bond in CO2 and H2O is GREATER than energy required to break the bonds in the

alkanes and O2

CH4 + O2 CO2 + H2O Hc = negative value

o Exothermic is an evidence for the formation of strong bond.

o Endothermic is an evidence for the formation of weaker bond.

Hess Law

• Hess law states that the enthalpy change in a chemical reaction is the same whether

the change is brought about in one stage or through intermediate stages. • It is used to determine enthalpy change that cannot be found by direct experiment.

Dr. Catherine Tan

10


Atomic Structure and the Periodic Table Relative Atomic Mass, Ar

• Atoms are too small to be weighed, therefore isotopes carbon-12 (12C) has been assigned

a mass of exactly 12 atomic mass unit (a.m.u) for comparison purpose.

• The mass of one atom of an element can be found by comparing with carbon-12 atom and

the mass obtained is known as Relative Atomic Mass.

Relative Molecular Mass, Mr

• It is defined as:

Dr. Catherine Tan

11


Mass Spectrometer

Mass spectrometer can be used to determine:

• The relative atomic mass, Ar of an atom,

• The relative molecular mass, Mr of a molecule,

• The relative abundance of an isotope in a sample of an element.

• The structure or identity of a compound.

Dr. Catherine Tan

12


Mass Spectrum Of Element Chlorine

Chlorine has two isotopes, 35-Cl and 37-Cl, in the approximate ratio of 3 atoms of 35-Cl to 1 atom

of 37-Cl

Ar Cl = (3/4 x 35 + 1/4 x 37)

= 35.5

Moles & Avogadro Constant

o Avogadro constant, L is defined as the number of carbon atoms in exactly 12 a.m.u of

carbon-12 which is 6.02 x 1023.

o Mole is defined as the amount of substance containing Avogadro’s number of particles of

that substance.

o For instance,

1 mole of carbon-12 = 6.02 x 1023 atoms

1 mole of H2 = 6.02 x 1023 molecules

1 mole CO2 = 6.02 x 1023 molecules

o 1 mole of any substance has a mass equal to its Ar or Mr

Dr. Catherine Tan

13


Atoms

o Atoms are made up of three fundamental: protons, neutrons and electrons

o Protons and neutrons are found in the nucleus of an atom

o Nucleus provide nearly all the mass of an atom

o Nucleon number = neutron number + proton number

o Neutral atom: electron = proton

o Anion: electron > proton

o Cation: proton > electron

Atomic Orbital

o Orbital: a region where the electrons occupy around the nucleus of an atom

o Atomic orbital: A region where there is the greatest possibility of finding a particular

electron in a free atom

o Electrons occupy 4 types of orbital: s, p, d and f

o S orbital: spherical in shape

o P orbital: px, py and pz dumb-bell shaped and are arranged along x, y and z axes in

space

Dr. Catherine Tan

14


Electron Shells and Sub-shells

o Electron shell (or principle shell, n) of an atom contains a group of orbitals which are

same distance from the nucleus

o Sub-shell is a group of orbital with same energy level but different orientation in

space

o The number of sub-shells in a principle shell is the same as the principle quantum

number of the principle shell

Energy Level of Orbital

Dr. Catherine Tan

15


Electronic Configurations

o The electronic configuration of an atom can be determined using 3 rules:

- The Aufbau Principle: electrons must occupy available orbital of lower energy first

before they fill orbital with higher energy

- Pauli Exclusion Principle: each orbital can occupy by 2 opposite spin only. Paired

electrons can only stable when they spin opposite direction so that the magnetic

attraction which result from their opposite spin can counterbalance the electrical

repulsion that result from their identical charges. The spins of electrons are

represented by and

- Hund’s rule: In a set of orbital of equivalent energy, electrons tend to occupy the

orbital singly first before pairing. For instance,

Dr. Catherine Tan

16


o Energy level of 4s orbital is lower than that of 3d orbital, so electrons would occupy 4s

orbital first before filling 3d orbital.

o However, once electrons filled into 3d orbital, the energy level reversed.

o For example, the electronic configuration of scandium (Z= 21) is 1s2 2s2 2p6 3s2 3p6 3d1

4s2 and NOT 1s2 2s2 2p6 3s2 3p6 4s2 3d1

o There is an exception in Chromium (Z =24) and Copper (Z = 29)

Dr. Catherine Tan

17


o An atom is in the ground state when all the electrons are in the lowest available energy

level (Usually at room temperature).

o When it is in excited state, one or more electrons absorb energy and promoted to a higher

energy level.

Ground state: 1s2 2s2 2px1 2py

1 2pz0

Excited state: 1s2 2s1 2px1 2py

1 2pz1

o Atoms that have the same number of electrons are known as Isoelectronic species

o Using Noble gas ‘Core’ in writing electronic configuration.

o For example,

Carbon 1s2 2s2 2p2 can be written as [He] 2s2 2p2

Chromium 1s2 2s2 2p6 3s2 3p6 3d5 4s1 can be written as [Ar] 3d54s1

Copper 1s2 2s2 2p6 3s2 3p6 3d10 4s1 can be written as [Ar] 3d10 4s1

Ionisation Energy

o It is defined as the amount of energy required to remove one electron from each atom in a

mole of gaseous atoms producing one mole gaseous cations under standard state (1atm,

25C)

o Ionisation energy is a positive value since energy is absorbed to remove an electron.

o Successive ionisation energies of an element increase with removal of each electron

because the remaining electrons are more tightly bonded by the positive charged in the

nucleus.

o No of ionisation energy = Atomic numbers

Dr. Catherine Tan

18


Factors affecting ionisation energy

o The distance between the nucleus and the electron

The attraction between the nucleus and the electrons decreases with increasing

distance between them and thus, the larger the size of an atom, the lower the

ionisation energy

o The nuclear charge

The higher the nuclear charge, the stronger the attraction between the nucleus

and the electrons and hence causes the ionisation energy to increase

o The screening effect

When the number of inner shells that filled with electrons increases, the

valence electrons are more shielded from the attraction of the nucleus and so

lower the ionisation energy

The repulsion between the inner filled electron shells will causes the size of an

atom increase, therefore, it will decrease the ionisation energy

Trends of ionisation energy across a period

o Second period

o Third period

There is an increase in the 1st ionisation energy moving from left to right of Period 2

and Period 3. This is because, moving from left to right, the atomic size decreases but

nuclear charge increases, hence the electrons are more tightly bound to the nucleus.

Dr. Catherine Tan

19


Exceptions moving from left to right of 2nd and 3rd period

o From the graph, 1st ionisation energy of Al is lower than that of Mg. This is because less

energy is needed to remove an electron from 3p orbital in Al as compared to 3s orbital in

Mg.

Mg: 1s2 2s2 2p6 3s2 Al: 1s2 2s2 2p6 3s2 3p1

o 1st ionisation of energy B is lower than Be

Be: 1s2 2s2 B: 1s2 2s2 2p1

o 1st ionisation of S lower than P

P: 1s2 2s2 2p6 3s2 3px1 3py

1 3pz1

S: 1s2 2s2 2p6 3s2 3px2 3py

1 3pz1

o 1st ionisation of O lower N

N: 1s2 2s2 2px1 2py

1 2pz1

O: 1s2 2s2 2px2 2py

1 2pz1

o 1st ionisation of Na is lower than Ne

Ne: 1s2 2s2 2p6 Na: 1s2 2s2 2p6 3s1

Dr. Catherine Tan

20


Trends of ionisation energy down a group

o When going down a group, the size of the atoms increases while the nuclear charge

decreases, hence the attraction between the nucleus and the electrons decrease and so

lower ionisation energy.

Successive Ionisation Energy

o Total number of electrons in an atom – this is equal to the separate number of ionisation

energies possessed by the atom

o Number of quantum shells occupied and the number of electrons in each – these deduced

by plotting successive ionisation energies against the order of removing the electrons

from the atom.

o Number of sub-shells occupied and the number of electrons in each – deduced by plotting

successive energies in a quantum shell against the order of removal of electrons.

o For example, Magnesium (1s2 2s2 2p6 3s2) is in group 2 of the periodic table and has

successive ionization energies

o Big jump occurs after second ionization energy. It means that there are 2 electrons

which are relatively easy to remove (3s2 electrons), while third one is much more

difficult because it comes from an inner level, closer to nucleus and with less

screening

Dr. Catherine Tan

21


o If we plot graphs of successive ionization energies for a particular element, we can see

fluctuations in it cause by different electrons being removed

o Not only we can see big jumps in ionization energy when an electron comes from an

inner level, but we can also see minor fluctuations within a level depending on whether

electron is coming from an s or a p orbital, and even whether it is paired or unpaired in

that orbital

o For example, graph plots first eight ionization energies of chlorine

Chlorine (17) 1s2 2s2 2p6 3s2 3p5

Dr. Catherine Tan

22


Bonding Ionic Bonding

• It is a strong electrostatic attraction force between oppositely charged ions.

• Metal with low ionisation energy (IE) tends to lose it valence electron to form a

positively charged ion (cation).

• Non-metal with high electron affinity (EA) tends to receive an electron to form a

negatively charged ion (anion).

Dot-and-Cross diagrams

• NaCl

• MgO

Dr. Catherine Tan

23


Lattice Structure of Sodium Chloride

Physical Properties of Ionic Compound

• Ionic compounds have high melting and boiling point because of the strong electrostatic

force between opposite charged ions

• Many ionic compounds are soluble in water although not all. It depends on whether

there are big enough attractions between water molecules and ions to overcome

attractions between ions themselves

• Ionic compound is insoluble in organic solvents

• Ionic compound is poor conductor in solid state because there are no ions which are

free to move. They only can conduct electricity in molten/liquid states through

electrolysis

Uses of magnesium oxide

• Magnesium oxide is a poor conductor, it is used as an electrical insulator in heating

elements and industrial cables

• Magnesium oxide is used for production of ceramics, transparent glass and crockery

• Magnesium oxide has high melting and boiling point makes it as a basic refractory

material for furnace lining. The term refractory refers to the quality of a material to retain

its strength at high temperatures

• A furnace is a device used for heating eg. extraction of metal from ore, combustion fuel

etc

Covalent Bonding

• It is a strong attraction force formed by sharing of electrons between two non-metallic

atoms.

• It is usually formed when the electronegativity between the two atoms are small.

• It can be a single bond, double bond or triple bond depends on the number of share pair

of electrons.

Dr. Catherine Tan

24


• The strength of a covalent bond is depends on the magnitude of the attraction between

bonded nuclei and shared pair of electrons and it is measured in terms of Bond energy

and Bond length.

Dot-and-Cross diagrams

• Hydrogen

• Oxygen

Exceptions in Octet Rule

• Compound with more than 8 electrons in the outer shell per atom.

Dr. Catherine Tan

25


• Compound with less than 8 electrons in the outer shell per atom.

• Only period 3 elements and beyond can expand its octet to accommodate more than 8

electrons.

• Period 2 elements can only accommodate a maximum of 8 electrons in its outer shell.

Dative Bond

• A dative bond formed when both electrons in a covalent bond are supplied by one of

the bonded atom instead of sharing between two bonded atoms.

• In order to form a dative bond, the donor atoms must have lone pair of electrons in its

outer shell while the acceptors must have vacant orbital in its outer shell.

• Lone pair of electrons is a pair of non-bonding electrons under the control of one atom.

Examples

• NH4

• Al2Cl6

Dr. Catherine Tan

26


Physical Properties of Covalent Compound

• It has low melting and boiling points such as methane and ammonia.

• It is the weak intermolecular forces that responsible for the physical properties of

covalent compound not the strong covalent bond that exists within a molecule.

Sigma and Pi Bonds

• The study of covalent bond formation in terms of atomic orbital overlap is known as

Valence Bond Theory.

• It involves two types of covalent bond, σ bond and π bond.

• Sigma Bond – Formed when two orbitals from two atoms overlap end-on or

known as head-to-head overlapping.

Pi Bond – Formed when two p orbitals of two atoms overlap sideways or known as

side-to-side overlapping.

π bond is weaker than σ bond since the overlapping of electron clouds in π is lesser.

Dr. Catherine Tan

27


• In a covalent compound, single bond is a σ bond whereas double bond consists of 1 σ

bond and 1 π bond.

• Triple Bond: two σ bond whereas double bond consists of 1 σ bond and 1 π bond.

Dr. Catherine Tan

28


Shape of Molecules

• The electron pairs around the central atom of a molecule dictate the shape of the

molecule.

• Steps in determining the shape of molecule:

1. Decide which atom is the center of a molecule.

2. Determine the no. of electron pairs around the central atom.

i. Look up valence electrons in central atom

ii. Add one electron for each atom joined to the central atom.

iii. Add electron if particle is negatively charged or subtract electron if the

particle is positively charged.

3. Determine the no. of bond pairs and lone pairs.

Valance Shell Electron Pair Repulsion (VSEPR) stated that:

– The electron pairs (either bond pair or lone pair) repel each other and move as far

apart as possible.

lone-pair lone-pair repulsion > lone-pair bond-pair repulsion> bond-pair bond-pair repulsion

• The repulsion between electron pair increased by an increase in electronegativity of the

central atom.

Dr. Catherine Tan

29


Molecule Electron dot diagram No. of bonding

pair

No. of lone pair Shape Geometry

BeCl2 Linear

H2O Bent

Dr. Catherine Tan

30


BF3 Trigonal planar

NH3 Pyramidal

Dr. Catherine Tan

31


CH4

Tetrahedral

PF5 Trigonal bipyramid

Dr. Catherine Tan

32


SF6 Octahedral

Dr. Catherine Tan

33


No. of

bonding pair

No. of

lone pair

Shape of molecule Example

2 0 Linear BeCl2

3 0 Trigonal planar BF3

4 0 Tetrahedral CH4

3 1 Trigonal pyramidal NH3

2 2 Bent H2O

5 0 Trigonal bipyramid PF5

6 0 Octahedral SF6

Dr. Catherine Tan

34


Bond Energy and Bond Length

• Bond energy is defined as the standard enthalpy change for breaking the bond in 1 mol

gaseous molecules.

• Bond length is the distance between the nuclei of two bonded atoms. The shorter the bond

length, the higher it is the bond energy.

Bond Polarity

• Covalent bonds may have some ionic character which resulting in a polar covalent bond.

• A polar covalent bond is formed between two atoms of different electronegativities.

• For example, Hydrogen fluoride consists of a polar covalent bond. This is due to the high

electronegativity of fluorine atom which tends to exert a stronger attraction on the

bonding electrons as compared to hydrogen atom.

• This unequal sharing of electrons is known as polarisation and the covalent bond is said

to be polarised.

• Therefore, a molecule is said to be polar (or has dipole moment) when its bonds are

polarised and it is not symmetrical.

• Ionic bond also contain some covalent character due to polarisation (cation attracts the

negative charge of the anion.

Dr. Catherine Tan

35


Intermolecular Forces

• High degree of covalency in ionic bond exist when,

– The cation is small

– The anion is large

– The charge on both ions is large

• There are two types of intermolecular forces:

– Van der Waals’ forces

– Hydrogen bonding

– Van der Waals’ forces are forces of attraction can be divided into two,

– Dipole-dipole forces

– Temporary dipole-induced dipole forces

• Dipole-dipole forces exist between polar molecules.

• The positive end of the dipole of one molecule will attract the negative end of the dipole

of another molecule.

• As for non-polar molecules such as oxygen and nitrogen, it is suggested that there are

force of attraction between molecules since they can be liquefied and solidified.

• Temporary dipole-induced dipole attraction is due to the temporary fluctuations in the

electron density of a molecule.

Factor affecting the strength of Van der Waals’ forces

• No. of electrons - The greater the no. of electrons in the molecules, the stronger is the van

der Waals’ forces of attraction.

• Shape of molecule - For instance, strength of van der Waals’ is reduced when there is a

branching because smaller surface area of contact for van der Waals’ forces.

• Hydrogen bond can exists between molecules that contain a hydrogen atom covalently

bonded to a small and highly electronegative atom, for instance fluorine, oxygen and

nitrogen. This is also known as intermolecular hydrogen bonding

Dr. Catherine Tan

36


• The highly electronegative atom will attract electron density towards itself and causes a

dipole moment.

• Hence, the hydrogen atom of one molecule will possess a partial positive charge and will

be attracted to the fluorine, oxygen or nitrogen atom of another molecule that carries a

partial negative charge.

Intermolecular & Intramolecular Hydrogen Bonds

• In 2-nitrophenol, there is possibility of an intramolecular hydrogen bond forming

between hydrogen atom of hydroxyl group (-OH) and oxygen atom of nitro group (-NO2)

Dr. Catherine Tan

37


• However, in 4-nitrophenol, the functional groups are on opposite sides of benzene ring.

Hydroxyl group is too far away from nitro group to form intramolecular hydrogen bond.

So, intermolecular hydrogen bonds are formed between molecules

• Intermolecular bonding in 4-nitrophenol is far stronger than in 2-nitrophenol, so boiling

point 4-nitrophenol is higher than 2-nitrophenol

Physical Properties of Hydrogen Bonding

• Hydrogen bonding is responsible for the high boiling point of water and low density of

ice.

• Water might exist as gas at room temperature without existence of hydrogen bonding.

• In the presence of hydrogen bonding, ice has an open structure which account for the

lower density of ice as compared to water.

Dr. Catherine Tan

38


Giant Covalent Molecules

Diamond

• Carbon has an electronic arrangement of 1s2 2s2 2p2. In diamond, each carbon shares

electrons with four other carbon atoms - forming four single bonds.

The physical properties of diamond

• Has a very high melting point (almost 4000°C). Very strong carbon-carbon covalent

bonds have to be broken throughout the structure before melting occurs.

• Is very hard. This is again due to the need to break very strong covalent bonds operating

in 3-dimensions.

• Doesn’t conduct electricity. All the electrons are held tightly between the atoms, and

aren't free to move.

• Is insoluble in water and organic solvents. There are no possible attractions which could

occur between solvent molecules and carbon atoms which could outweigh the attractions

between the covalently bound carbon atoms.

Graphite

• Graphite has a layer structure.

Dr. Catherine Tan

39


• Each carbon atom in graphite uses three of its electrons to form simple bonds to its three

close neighbours. That leaves a fourth electron in the bonding level. These "spare"

electrons in each carbon atom become delocalised over the whole of the sheet of atoms

in one layer. They are no longer associated directly with any particular atom or pair of

atoms, but are free to wander throughout the whole sheet.

The physical properties of graphite

• It has a high melting point, similar to that of diamond. In order to melt graphite, it isn't

enough to loosen one sheet from another. You have to break the covalent bonding

throughout the whole structure.

• It has a soft, slippery feel, and is used in pencils and as a dry lubricant for things like

locks. You can think of graphite rather like a pack of cards - each card is strong, but the

cards will slide over each other, or even fall off the pack altogether. When you use a

pencil, sheets are rubbed off and stick to the paper.

• It has a lower density than diamond. This is because of the relatively large amount of

space that is "wasted" between the sheets.

• Insoluble in water and organic solvents - for the same reason that diamond is insoluble.

Attractions between solvent molecules and carbon atoms will never be strong enough to

overcome the strong covalent bonds in graphite.

• Conducts electricity. The delocalised electrons are free to move throughout the sheets. If

a piece of graphite is connected into a circuit, electrons can fall off one end of the sheet

and be replaced with new ones at the other end.

Silicon Dioxide

• Silicon dioxide is also known as silicon (IV) oxide.

• Crystalline silicon has the same structure as diamond. To turn it into silicon dioxide, all

you need to do is to modify the silicon structure by including some oxygen atoms.

The physical properties of silicon dioxide

• It has a high melting point - varying depending on what the particular structure is

(remember that the structure given is only one of three possible structures), but around

Dr. Catherine Tan

40


1700°C. Very strong silicon-oxygen covalent bonds have to be broken throughout the

structure before melting occurs.

• It is hard. This is due to the need to break the very strong covalent bonds.

• It doesn't conduct electricity. There aren't any delocalised electrons. All the electrons are

held tightly between the

Metallic Bonding

• It is defined as the electrostatic attraction between the positively charged metal ions and

the ‘cloud’ of delocalised electrons

The physical properties of metal

• Metals can be deformed since the electron ‘cloud’ prevent the repulsion among the

cations.

• Since it electrons are free to move throughout the metal piece, hence metal is a good

electrical and thermal conductors.

• Owing to these typical properties, metal such as aluminum, copper can be used for

packaging, drawn into wires, and their alloys are strong enough for usage in aeroplane,

car, train and building.

Dr. Catherine Tan

41


Introductory Organic Chemistry

Skeletal Formula

Dr. Catherine Tan

42


Nomenclature

o The naming of organic compounds follows the IUPAC (International Union of Pure and

Applied Chemistry).

o Every name consists of 3 parts:

The 1st part indicates the number of carbon atoms in the longest continuous chain

(parent chain).

1 carbon = meth- 5 carbon = pent-

2 carbon = eth- 6 carbon = hex-

3 carbon = prop- 7 carbon = hept-

4 carbon = but- 8 carbon = oct-

The 2nd part indicates the linking or bonding in the chain.

-an- means all single bond in the carbon chain

-en- means a double bond in the carbon chain

-yn- means a triple bond in the carbon chain

The 3rd part indicates what functional group is joined to the chain.

-e means only H joined to the carbon chain (alkanes/alkenes/arenes)

-ol means a -OH group in the carbon chain (alcohol)

-amine means a –NH2 group in the carbon chain ( amines)

-al means a C=O group on the end of the carbon chain (aldehydes)

-one means a C=O group in the carbon chain but not at the end (ketones)

-oic acid means a CO2H group in the carbon chain (carboxylic acid)

Dr. Catherine Tan

43


Functional Group

o It is a group of atoms within a compound whose reactions dominate the chemistry of the

molecule and so gives the characteristic properties.

o Homologous series is a series of compounds containing the same functional group and

adjacent members differ in their formula by a CH2 unit.

Dr. Catherine Tan

44


Homolytic and Heterolytic Fission

o Organic reaction involve breaking and formation of covalent bonds.

o Two ways in which a covalent bond can be break:

Homolytic fission: the breaking of a covalent bond such that one electron goes to each

of the atom forming free radicals.

Dr. Catherine Tan

45


Free radical is an atom or group of atoms with an unpaired electron formed from

the homolytic fission of a covalent bond and are very reactive

Heterolytic fission: the breaking of a covalent bond such that both the electrons go

to the same atom forming positive and negative ions.

A carbocation is a carbon species that carries a positive charge.

A carbanion is carbon species that carries a negative charge.

Nucleophiles & Electrophiles

o Nucleophiles are species which contain a lone pair of electrons and are attracted to regions of

positive charge or electron deficiency sites in a molecule.

o e.g. NH3, CN- OH-, Cl-, Br-, R-NH2, H2O.

o Electrophiles are electron-deficient species which can accept electrons and are attracted to

regions of negative charge or electron rich sites in a molecule.

e.g. H+, Br+, Cl+, NO2+, R+.

o If a nucleophile is being used, the reaction is called nucleophilic.

o If an electrophile is being used, the reaction is called electrophilic.

o If an atom or group has been added, the reaction is called an addition.

Dr. Catherine Tan

46


o If an atom or group is replaced by another, the reaction is called substitution.

o If atoms from two neighbouring carbons are lost, the reaction is called an elimination.

Organic Reactions

o Addition: Involves two molecules joining to form a single new molecule.

o e.g.

o Substitution: Involves replacing an atom ( or a group of atoms) by another atom (or a group

of atoms).

o e.g.

o Elimination: Involves the removal of a small molecule from a larger molecule.

o e.g.

Dr. Catherine Tan

47


o Oxidation: Involves addition of O atoms by the reaction with oxidising agent.

o e.g.

o Reduction: Involves addition of H atoms by reaction with reducing agent.

o e.g.

Isomerism

Isomerism occurs when two or more compounds have the same molecular formula but different

arrangement of the atoms in the molecules. These compounds are known as isomers

Dr. Catherine Tan

48


o Chain isomerism arises due to different arrangement of carbon atoms in a chain. The carbon

atoms may arranged in a straight chain or branched chain.

o e.g.

o Position isomerism arises due to the different positions of a functional group in carbon chain.

o e.g.

o Functional group isomerism arises due to different functional groups.

o e.g.

o Geometrical Isomerism (cis-trans)

• Geometrical isomerism arises when rotation about a bond is restricted.

• It is common in compounds with C=C bonds where rotation is restricted sue to the

presence of π bond.

• The cis-isomer has two groups attached to carbon atom on the same side of the double

bond whereas trans-isomer has two groups on the opposite side of the double bond.

Dr. Catherine Tan

49


• Cis-trans isomerism cannot exist if either carbon carries two identical groups.

• e.g.

Cis-trans isomers have similar chemical properties. They react with same reagents but at

different rates

• Cis-trans isomers have different physical properties.

– Cis-isomer has a higher boiling point because of its higher polarity.

– Cis-isomer has lower melting point because of its lower symmetry.

– e.g.

Dr. Catherine Tan

50


Dr. Catherine Tan

51


Alkanes o Alkanes are saturated hydrocarbons in which the carbon atoms are joined by single covalent

bond only.

o They form a homologous series with general formula CnH2n+2.

Physical Properties of Alkanes

o Alkanes are soluble in non-polar solvent such CCl4, ether and benzene.

o The density of liquid alkanes increases slightly with increased size of the molecules due to

increasing intermolecular van der Waals’ forces which causes alkanes to be more compact in

the condensed liquid state.

o The boiling point of the straight chain alkanes increases steadily with increased size of

molecules. This is due to the increased intermolecular forces as the number of electrons in

the molecule increases.

o Branching increases the volatility and reduces the density of the alkanes.

o With branching, the molecules become more spherical in shape and they pack together less

closely, resulting in lower density and smaller surface areas of contact for van der Waals’

forces. Hence, boiling point and melting point decreases as the strength of van der Waals’

forces decreases.

Dr. Catherine Tan

52


Combustion of Alkanes

o Alkanes burn exothermically in excess oxygen to form CO2 and H2O. (complete combustion)

o The ease of burning and their exothermic reactions account for the use of many alkanes as

fuels.

o Alkanes only burn in gaseous state and those less volatile alkanes will burn less readily.

o Solid and liquid alkanes must be vaporised before they will burn.

o When alkanes burn in a limited supply of oxygen, CO and C will form. (incomplete

combustion)

o As a results of incomplete combustion, CO and oxides of nitrogen, NOx which are pollutants

to the environment.

o A catalytic converter is used to remove CO, NOx and unburnt hydrocarbons. These harmful

gases are converted into less harmful CO2, N2 and water vapour.

o In catalytic coverter, oxides of nitrogen are decomposed to oxygen and nitrogen or reduced

by CO (or unburnt hydrocarbons).

o CO and unburnt hydrocarbons are oxidised to CO2 and H2O.

Substitution Reaction of Alkanes

o Substitution by halogen is known as Halogenation.

– Reagent: Cl2 or Br2 in CCl4

– Condition: In the presence of UV light

– Product: Chloroalkanes or Bromoalkanes

– When methane reacts with chlorine in sunlight, the yellowish-green colour of

chlorine fades and steamy acidic fumes of HCl can be detected.

o This is a chain reaction; one or more chlorine atoms may replace hydrogen atoms, depending

on the amount of halogen and alkane present.

Dr. Catherine Tan

53


Free-Radical Substitution

Dr. Catherine Tan

54


Alkenes o Alkenes are unsaturated hydrocarbons with general formula CnH2n.

o Alkenes contain C=C bonds in their structures.

o If the alkenes molecule contain two double bonds, it is named as diene.

o e.g. CH2=CH-CH=CH2 buta-1,3-diene

CH2=CHCH2CH=CH2 penta-1,4-diene

Reactivity of Alkenes

o Alkenes are much more reactive than alkanes, with the charge clouds of the π bond being the

reactive site.

o The reaction of alkenes are mainly addition reaction, involving the π electrons in which the

unsaturated alkenes are converted into alkanes.

o e.g.

Dr. Catherine Tan

55


Addition of hydrogen:

Addition of steam:

Dr. Catherine Tan

56


Markovnikoff’s Rule

• In addition of H-X to a C=C bond in an unsymmetrical alkenes, the H atom attaches itself

to the carbon atom that already holds the greater number of hydrogen atoms.

• e.g. CH2=CHCH3 + H2O CH3-CH(OH)CH3

• This is because the reaction involves the formation of an intermediate carbocation, the

stability of which decreases in the order:

• Tertiary (3o) carbocation is most stable because it has three electron-donating alkyl (R)

groups which neutralise the positive charge more than of the secondary and (2o) and

primary (1o)

Addition of hydrogen halides:

• Reagent: HX (where X = Cl, Br or I)

• Condition: Room temperature

• Product: Halogenoalkanes

– The reaction again involves the formation of an intermediate carbocation.

– The addition of HX to unsymmetrical alkenes follows Markovnikoff’s rule.

– e.g. CH2=CHCH3 + HCl CH3CH(Cl)CH3

– The rate of this reaction decrease in the order: HI> HBr> HCl> HF

– Addition of halogen:

– Reagent: X2 (where X = Cl, Br or I)

– Condition: Room temperature (in the absence of UV light).

Dr. Catherine Tan

57


– e.g. When ethene is bubbled into Br2 in CCl4 at room temperature (no

UV light), the reddish-brown colour is rapidly

decolourised.

–

– e.g. When ethene is bubbled into aqueous Br2 at room temperature (no

UV light), the reddish-brown colour is rapidly

decolourised. Two products formed, 1,2- dibromoethane and 2-

bromoethanol.

– The reaction of alkenes with Br2 in CCl4 is a test for unsaturation.

–

Addition Polymerisation:

• An addition reaction is one in which two or more molecules join together to give a single

product. During the polymerisation of ethene, thousands of ethene molecules join

together to make poly(ethene) - commonly called polythene.

n CH2=CH2 -[-CH2CH2-]-n

ethene poly(ethene)

• The number of molecules joining up is very variable, but is in the region of 2000 to

20000.

• Poly(ethene) are non biodegradable makes them hard to dispose and as a result, they can

act as breeding places for many of the disease germs which, sooner than later cause an

epidemic in the surrounding people.

• Environmentally unfriendly considering the time taken for their decomposition. As a

result of this time spun they can cause further problems like blocking water penetration

into the soil which in turn affects food growth and development

unit 1: the core principles of chemistry - … · unit 1: the core principles of chemistry ... o...

Documents