Electron density maps
Bonds consist of electrons shared between molecules, which can b portrayed by drawing maps showing the contours of the electron rings.
Electron Sharing in covalent molecules
(Dot and cross diagram)
Pairs of electrons in an atom tent to get as far away from each other as they can. This results in a tetrahedral distribution when there are four electron pairs.
Non-bonding pairs of electrons in an atom, (pairs of electrons not shared) repel more strongly than shared pairs do. These are lone pairs.
Multiple bonds
= where more than one electron from each atoms is shared covalently. E.g. Ethene
This is the double bond formed by 2 pairs of electrons covalently bonding.
This double bond is referred to as a pi bond (π) where the electrons form two electron clouds. (The sigma (σ) bond is till present, the single line in the diagram below). The π is not equivalent to a σ bond as it is much weaker having a lower energy.
Shapes of Molecules
Molecules are 3D
Bond angles are determined by repulsive forces between electrons, bonded and lone pairs.
E.g. Methane CH4
2D suggests that the bond angles are 90ْ , but,
The shape is determined by VSEPR (Valence Shell Electron Pair Repulsion) theory. With four electron pairs around a central atom, the methane above is tetrahedral; the bond angle for a tetrahedral molecule is 109.5.
Overall as the number of bonded pairs of electrons increases the bond angle increases. Also the less pairs of lone electrons there are the greater the bond angle.
Covalent Giant Structures
These are 3 dimensional networks of atoms bonded with covalent bonds.
Carbon can form giant covalent structures, diamond and graphite.
Each atom forms four covalent bonds with another carbon atom which arrange into a 3D lattice.
Properties
Not soluble in water
Particles have no charge
Polar water molecules not attracted to them
Don’t conduct electricity, no free ions or electrons to carry charge (except graphite where electrons are free to move between layers of lattice structure as only 3 electrons covalently bond. The electrons provide a weak attractive force between the layers, (delocalized electrons). Therefore graphite can conduct electricity and is the only non-metal to do so. See student book for diagrams of giant structures (p. 145).
Dative Covalency
= where a pair of covalently shared electrons have come from one atom.
When draw with bonds the bond is normally draw as an arrow towards the atom which has had the electrons donated to it.
This type of covalency also occurs in other molecules such as, carbon monoxide (CO), Ammonium ion (NH4+) and in Nitric acid, (NHO3).
Electronegativity
= the power of an atom in a molecule to attract electrons to itself. This unequal sharing of electrons is called polarisation.
These are the trends of electronegativity in the periodic table.
The attraction is exerted by the nuclei. When the forces exerted by the two nuclei are to different degrees which displaces electrons towards on atom.
The Chlorine atom in this molecule is more electronegative than the H atom. δ+ - δ-
This is an example of how completely ionic and covalent molecules are extreme types and molecules occur over a range of intermediate types. In conclusion, molecules can be partially ionic and partially covalent.
Where the centers of positive and negative charge are not coinciding then a permanent dipole occurs.
Bond Energies
When bonds are formed, energy is released = exothermic.
When bonds are broken, energy is required = endothermic.
The amount of energy associated with a bond is fixed, called bond energy orenthalpy.
E.g. Br – Br = -192.9 KJ mol-1
Br – Br → 2Br ∆H = + 192.9 KJ mol-1
2Br → Br – Br ∆H = - 192.9 KJ mol-1
The standard enthalpy change of combustion of a substance, symbol ΔHөc is defined as the enthalpy change that occurs when one mole of the substance undergoes complete combustion under standard conditions
The overall energy change in going from elements to combustion products must be the same whatever route. So Route 3 = 1 + 2
-965.1 KJmol-1
Energy Change Using Bond Energies
E.g. CH4 + 2O2 → CO2 + 2H2O
Bonds Broken Bonds Formed
4 E(C – H) = 4 x 435 2 E(C = O) = 2 x 805
2 E(O = O) = 2 x 498.3 4 E(O = H) = 4 x 464
= 2736.6 = 3466
Energy difference = bonds formed – bonds broken
= 3466 – 2736.6
∆Hө = -729.4 KJ mol-1
Bond Lengths
As Bond length decrease the bond energies increase
If the force of attraction between atoms is greater they will be close together giving a shorter bond length and greater bond energy.
This larger force of attraction is caused by:
- The number of electrons in the bond (single, double, triple)
- Large dipole, larger the dipole , the stronger the bond
Bond Energies Rates of reactions
In order to react, molecules must collide, which must have enough energy to break the bonds in the molecule and have the correct orientation when colliding.
The energy required for this to happen is known as the activation energy.
The activation energy can be lowered by a catalyst to speed up a reaction; the catalyst provides an alternative pathway to a lower EA
Dynamic Equilibrium
Reversible reactions (↔)
E.g. H2(g) + I2(g) ↔ 2HI(g)
Different conditions favour both forward and backward reactions.
E.g. a higher temperature favour the forward reaction is the reaction is endothermic, whereas a lower temperature favours the backward reaction, where it is exothermic.
In a closed system where no matter can leave of enter, the reaction reaches equilibrium where the two reactions are occurring at the same rate.
The position of equilibrium can be changed by changing the conditions.
Le Chatlier’s Priciple: “When a system in equilibrium is subjected to a change, the processes which take place are such to tend to counteract the change.”
This means if a change takes place the reaction will change to undo what has been done.
For example, if the conc. of reactants is increased the equilibrium will move to use the reactants up
If you increase the temperature then the equilibrium will change to cool it down.
Effect of Conditions
- Lower Temp favours exothermic
- Higher Temp favours exothermic
- Higher pressure favours fewer gas molecules
- Lower pressure favours more gas molecules
Topic 8 Organic Chemistry Hydrocarbons
Photochemical reactions between Halogens and alkanes
Alkanes only react with halogens when exposed to U.V. light – Photochemical reaction.
C6H14(l) + Br2(aq) → C6H13Br1(l) + HBr(g)
This is a substitution reaction, The H atom is removed and replaced with a Br atom. This is normally prevented from happening by the high bond energy (C-H).
But with U.V. light this can occur.
-
The U.V. light breaks the halogen molecule into free radicals (an atom with unpaired electrons) by homolytic fission.
- The highly reactive free radicals attack the alkane
Cl∙ + CH4 → H∙ + CH3Cl
Both these reactions produce more fee radicals
Cl∙ + CH4 → CH3∙ + HCl
- The reaction is terminated when two free radicals react with each other.
Cl∙ + Cl∙ → Cl2
CH3 ∙ + Cl ∙ → CH3Cl
CH3 ∙ + CH3 ∙ → CH3-CH3
Overall, C6H14(l) + Cl2(g) → C6H13Cl(l) + HCl(g)
Cl2 → 2Cl ∙ INITIATION
CH6H14 + Cl∙ → HCl + C6H13∙
Cl2 + C6H13 → C6H13Cl + Cl∙ PROPAGATION
Cl∙ + Cl∙ → Cl2
C6H13∙ + Cl∙ + → C6H13Cl
C6H13∙ + C6H13∙ → C12H26 TERMINATION
INITIATION = The production of free radicals
PROPAGATION = The products formed when free radicals attack alkanes/ free radical alkanes attack halogens, resulting in more free raicals
TERMINATION = When two free radicals combine
The Alkenes
Unsaturated hydrocarbons, contain multiple bonds, e.g. C=C
Alkenes have double bonds made from sigma and pi bonds.
Obtained from cracking alkanes.
Catalytic Cracking
Alkenes
Limonene is an alkene which can be extracted by steam distillation, for diagram see p. 195 in students book.(Unlike book, don’t leave gaps between quick fit)
Geometric Isomerism
(Isomer – same molecular formula; different structural formula)
trans - 1,2 – dichloroethene cis – 1,2 - dichloroethene
chlorines diagonally opposite chlorines on same side
This occurs as double bond cannot rotate.
Positional Isomerism
The fumctional group can be in different places on the chain (propan–1–ol, propan-2-ol)
This is the same with alkenes, the position of the double bond:
- C-C-C=C but-1-ene
- C-C=C-C but-2-ene
- C=C-C=C but-1,3-diene
Addition reactions of alkenes
When alkenes react with halogens, the halogen molecule is added directly into the molecule across the double bond.
-
Br2 approaches the ethane
-
The double bond repels the electrons in the Br2 resulting in an induced dipole.
-
The δ+ end of the Br2 is attracted to the π bond in the ethene. The π bond and Br-Br bonds break
- The new bonds form.
Topic 9 Intermolecular Forces
= forces between molecules not in them.
Evidence for states of matter
Gas → Liquid → Solid
In order for molecules to exist as solids or liquids there must be forces between the molecules holding them together.
Three basic types of forces;
Van der Waals forces – between any molecule
Dipole – dipole attractions – between polar molecules
Hydrogen bonding – between certain dipole molecules
Van der Waals forces (temporary dipole – dipole)
- weakest intermolecular attraction
- occurs in all molecules
Happens because electrons in the molecules move about and at certain times this can create a temporary dipole which means molecules can attract to each other.
Factors which can affect the strength of the forces
a)Molecular size/shape
Larger molecules have larger forces, greater surface area.
Branching reduces attraction
The molecules are closer when not branched meaning a greater attraction, whereas when a molecule is branched then the dipoles are further away reducing attraction.
The longer distances mean that the attraction is weaker.
b) Polarisation
How easily an electron cloud is distorted
Larger electron cloud, more polarisable. Larger electron cloud is more polarisable
Dipole – dipole forces
Stronger than V d W forces, due to being a permanent dipole.
Occurs in polar molecules (where electron clouds are already distorted due to delocalized electrons or higher electron density, due toelectronegativity)
E.g. Van der Waals and dipole – dipole
Increase in dipole moment =
Increase in attraction =
Increase in boiling point.
Hydrogen Bonding
Strongest intermolecular force.
Occurs in molecules with H bonded to O, N, or F, most electronegative bonding.
No electrons in H so becomes a large +ve dipole so highly attractive.
- Not a bond but a strong dipole – dipole attraction.
H bond is - - -.
E.g What predominates in these molecules?
SO2 dipole – dipole
CO2 – V d W
CH3CH2OH – H bonding
H2O – H bonding
H2S – H bonding
Dipole Moment
-Dipole moment measured in debyes (D)
5.0 D represents the transfer of 1 electron across a bond length (0.1 nm)
The dipole moment.
Evidence for H – Bonds
Group 0 boiling point increases down the group as there are more electrons increases which means the dipole will be larger so van der waals forces will be stronger. But this doesn’t occur in the groups V, VI and VII. In group IV it is the same pattern as there is in the hydrides (0) as they are non-polar so only V d W forces occur. In groups V, VI and VIII molecules are polar.
Therefore diple – dipole occur, but at top of groups are NH3, H2O, HF. These don’t follow trend as they have v. high boiling points, due to H – bonding, H2O has even higher boiling point due to having 2 H bonds per molecule.
Topic 10 Halogenoalkanes
Organic Compounds where as well as carbon bonded to H there are halogens (Cl, Br, I)
E.g.
chloromethane
bromoethane
1,1,2 – tribromo – 2 –chloro – 1,2, difluroethane
Carbons stay same number
Halogens in alphabetical order
Bond Energy
Bond energies are high therefore not very reactive
The bonds are strong due to the large dipole ()
The bonds get weaker down the group as the electronegativity of the elements decreases (see Electronegativity topic 7)
Due to the δ+ on the carbon these molecules are attached by negatively charged species (e.g. OH-) called nucleophiles resulting in nucleophilic substitution.
E.g. C2H5Br + KOH → C2H6O + KBr .
Primary, Secondary and tertiary halogenoalkanes, same as in alcohols, determined by the number of carbons are joined to the carbons which is attached to the halogen.
Reactions of Halogenoalkanes
Halgenoalkanes undergo nucleophilic substitution reactions, they are also made from alcohols by nucleophilic substitution.
Making Halogenoalkanes
CH3CH2OH CH3CH2Br + H2O
The OH- on the alcohol is called the leaving group and the Br- is a nucleophile called the attacking group.
The lone pair on the Br- is attracted to the δ+ on the carbon, caused by the OH group making a dipole.
The bond between the C-OH breaks and forms an OH- ion.Overall this is called a nucleophilic substitution reaction.
More reactions of halogenoalkanes
- Aqueous alkali
OH- acts as a nucleophile
CH3CCH3BrCH3 CH3CCH3OHCH3
2-bromo-2-methylpropane 2-methylpropan-2-ol
Br is the leaving group OH- the attacking group
This also occurs with H2O as the attacking group, but is very slow.
- With Ammonia
NH3 has a lone pair of electrons so acts as a nucleophile
CH3CH2Br CH3CH2NH2
Bromoethane ethylamine
- Alcoholic Alkali
In alcohol the OH- doesn’t act as a nucleophile, instead an elimination reaction occurs resulting in an alkene.
CH3CH2CH2Br CH3CH=CH2 + HBr
