Unit 1: The Core Principles of Chemistry - … · Unit 1: The Core Principles of Chemistry ... o...
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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.
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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
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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
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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
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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.
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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
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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.
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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.
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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.
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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:
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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.
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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
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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
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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
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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,
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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)
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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
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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.
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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
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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
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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
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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
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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.
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• 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.
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• 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
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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.
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• 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.
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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.
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Molecule Electron dot diagram No. of bonding
pair
No. of lone pair Shape Geometry
BeCl2 Linear
H2O Bent
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CH4
Tetrahedral
PF5 Trigonal bipyramid
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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
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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.
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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
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• 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)
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• 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.
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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.
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• 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
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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.
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Introductory Organic Chemistry
Skeletal Formula
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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)
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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.
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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.
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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.
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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.
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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
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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.
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• 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.
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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.
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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.
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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.
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Addition of hydrogen:
Addition of steam:
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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).
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– 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