Chemistry (A)

CFCs = used as refrigerants, propellants in aerosols, blowing agents for explanded plastics, cleaning solvents.

Good at jobs because of low reactivities, low boiling points, low toxicity and low stability.

However low stability causes problems as CFC have long life in troposphere. In the stratosphere, they undergo photodissociation - produces chlorine radicals which remove ozone.

In early 70s Sherry Rowland and Mario Molina predicted CFCs would damage ozone layer and were proved correct when hole in ozone was detected over Antartic using ultraviolet spectroscopy.

NASA satellites detected low ozone readings as anomalies. But aircraft few through stratosphere over Antarctica, measured concs of ClO and O3 - fall in O3 corresponded with rise of ClO.

Alternatives to CFCs = hydrochlorofluorocarbons (HCFCs) - H-C bonds broken down in trophosphere but are greenhouse gases that contribute to global warming - same with alkanes.

High-energy radiation from sun (visible and UV mostly) reaches Earth's surface - some frequencies absorbed. Earth's surface warmed - re-emits lower energy infrared radiation. Greenhouses gases (e.g. CH4) absorb some of this radiation - rest escapes into space.

Steady state reached - Earth radiates and absorbs energy at same rate. Absorption of infrared radiation by greenhouse gases causes atmospheric warming in two ways:

• Some infrared re-emited by molecules in all directions - back towards Earth and to space.
• Absorption increases vibrational energy of bonds in molecules so they vibrate more vigourously, energy is transferred to other molecules by collisions, increasing kinetic energy and raising the average temperature of the atmosphere.

CO2 and water = important greenhouse gases, absorb in two bands across Earth'sradiation spectrum.

Between the two bands there is a window where infrared radiation can escape without being absorbed - about 70% escapes through this range of frequencies.

NB: Global warming potential depends of how effeciently a gas absorbs infrared radiation and on its atmospheric lifetime.

Human activity = increasing atmospheric concentrations of natural greenhouse gases (e.g. CO2) and gases that are not present (CFCs). CFCs only present in small amount but have large global warming potential (GWP).

These gases absorb radiation in window where energy normally escapes through to space - leads to enhanced greenhouse effect.

To reduce global warming CO2 in atmosphere needs to be reduced - can be done by:

• Reducing consumption of fossil fuels
• Using alternative energy sources (wind, solar, tidal etc)
• Increasing photosynthesis.
• Burying or reacting carbon dioxide.

NB: Giant covently bonded structure = networks or giant structures.

Diamond:

• Carbon atoms, each joined tetrahedrally to four other carbon atoms by strong covalent bonds.
• Very strong C-C bonds and highly symmetrical network structure make diamond hardest naturally occuring structure.

Silicon (IV) oxide:

• Silicon atoms form four bonds, covalently to four oxygen atoms.
• Quartz = pure form of silicon (iv) oxide - etended network of SiO4 units, each silicon atom has a half share in four oxygen atoms.

Carbon dioxide = molecular structure w/ three atoms bonded in a linear arrangement O=C=O.

Has weak intermolecular bonds, so sublimes at a low temperature, intramolecular bonds are polar though, so it dissolves in water easily.

Silicon atoms = larger, have more electrons, unable to make double bonds so SiO2 molecules aren't formed.

Has giant network structure, requires lots of energy to overcome intramolecular bonds, so has high melting and boiling point and doesn't dissolve in water.

Atmosphere = divided into three sections: troposphere, stratosphere and ionosphere.

Troposphere = approx. 78% nitrogen and 21% oxygen, remaining is a mixture (mainly argon and CO2).

Human activity alters proportion of some naturally occuring gases (e.g. carbon dioxide). - become major pollutants.

Other pollutants e.g. CFCs (chlorofluorocarbons) = only occur as a result of human activity.

Concentration of gases = usually quoted in percent by volume, but with gases at less that 1% they are quoted in parts per million (ppm).

Molecules have certain energies associated with their behaviour.

In increasing order of energy these are:

• Translational energy - assoc. w/ molecule moving around as a whole.
• Rotational energy - assoc. w/ molecule rotating as a whole.
• Vibrational energy - assoc. w/ bonds vibrating within molecule.
• Electronic energy - assoc. w/ electrons moving from one energy level to another.

All these energies are quantised (exist at fixed levels). It is possible to move these energies from one fixed level to another = caused by radiation interacting with matter at specific frequencies.

In increasing order of frequency and decreasing order of wavelength:

Radiofrequency - Microwave - Infrared - Visible - UV - X-rays - Gamma rays

Different types of electromagnetic radiation = photons of different energy associated with them.

Electromagnetic radiation interacting with matter changes the energy associated with that matter.

Type of energy changes depends on type of radiation absorbed (in order of decreasing energy):

• Electronic change (caused by absorbing UV or visible).
• Vibrational change (caused by infrared radiation)
• Rotational change (caused by microwave radiation)

Energy of photons absorbed to cause these changes is related to the frequency of the radiation by the equation: E = hv, where E = energy, h = Planck's constant (6.63 x 10-34)  and v = frequency.

Electrons in molecules occupy definite energy levels.

One of three things happens when a molecule absorbs visible or UV radiation:

• Electrons = excited to a higher electronic energy level - will return to original states in time as they release the energy they absorbed.
• Chemical bonds can break - radicals form = photodissociation.
• Electron is ejected from a molecule, which then becomes ionised.

E.g. Cl2 + hv --> 

• Cl2 + e- = ionisation.
• Cl + Cl = dissociation.
• Cl2 = release of energy then returning to it's original state.

Energy absorbed: Ionisation > Dissociation > Release of energy (excitation).

In covalent bonds - pair of electrons shared between two atoms.

Bond breaks and electron pair is redistributed in one of two ways:

• Heterolytic fission = both electrons of pair go to just one of the atoms, forming ions.
• Homolytic fission = one of two electrons in the pair goes to each atom, so both atoms have one unpaired electron - radicals formed.

Radical = a species with one or more unpaired electrons.

If radical has two unpaired electrons it is called a biradical.

E.g. O2 = .O - O.

Radicals are very reactive because of unpaired electrons.

Radicals = highly reactive, undergo chain reactions.

Chain reactions divided into three stages:

• Initiation = there are no radicals at the beginning but radicals form at the end of the stage e.g. Cl:Cl -(hv)-> 2Cl.
• Propagation - radicals at the start of the stage and new radicals formed at the end e.g.           Cl. + H:H --> Cl-H + H. or H. + Cl:Cl --> H-Cl + Cl.
• Termination = reaction is terminated when two radicals collide e.g. H. + H. --> H-H

NB: Radical reactions are fast, often initiated by heat or light, normally occur in gas phase.

hv can represent energy.

A halogen can substitute a hydrogen in an alkane chain - via radical substitution mechanism.

Produces a halogenoalkane:

• Initiation = Homolytic fission occurs in presence of UV light: Cl2 -(hv)-> 2Cl.
• Propagation = CH4 + Cl. --> CH3. +HCl or CH3. + Cl2 --> CH3Cl + Cl.
• Termination = CH3. + CH3. --> C2H6 or CH3. + Cl. --> CH3Cl

When radicals form due to presence of light = photodissociation.

Dioxygen molecules in stratosphere (O2) can absorb UV radiation of right frequencey to split molecule apart = photodissociation - Oxygen atoms, which are radicals, are formed.

O2 -(hv)-> 2O.

Ozone (O3) formed when oxygen atom (radical) reacts rapidly with a dioxygen molecule:

O + O2 --> O3

Ozone is highly reactive - destroyed by reacting w/ radicals in stratosphere (if X is radical):

• X + O3 --> XO + O2
• XO + O --> O2 + X
• O + O2 --> O3

X from the second equation can continue to react as a reactant in the first equation - repeating shows how ozone is being lost from the system.

X is involved in the reaction but isn't used up - acts as a catalyst - example of a catalytic cycle, so one single chlorine atom can remove up to 1 million ozone molecules.

NB: Radical X (depletes ozone) could be:

• OH (hydroxyl radical) - formed from water.
• NO (nitrogen monoxide) - produced in internal combustion engines.
• Cl (chlorine radical) - produced in breakdown of CFCs from cleaning solvents, aerosols propellants and refridgerants.

In troposphere ground level ozone = irritant toxic gas which weakens immune system.

Some ozone is formed by sunlight acting on primary pollutants in photochemical smog - causes breathing problems for humans.

Ozone absorbs radiation between 10.1 and 14.0 x 10^14 Hz = ultraviolet region of electromagnetic spectrum - most damaging to the skin.

Ozone absorbs much of the UV radiation so damage to skin is reduced, but with less ozone the incidences in skin cancer are increasing.

Chemical reaction has a forward reaction AND a backward reaction - the backward reaction is significant then the reaction = reversible.

Dynamic equilibrium = when the rate of forward reaction is the same as the rate of the backward reaction - represented by <=> (normal symbol).

E.g. when carbon dioxide dissolves in water:

• CO2(g) <=> CO2(aq)
• H2O(l) <=> H2O(g)

These are physical changes.

This reversible reaction also occurs:

CO2(aq) + H2O(l) <=> HCO3-(aq) + H+(aq) (hydrogencarbonate ions and hydrogen ions form).

Concentrations of reactants and products remain unchanged once dynamic equilibrium is established but the reactions DO NOT STOP - they CONTINUE AT SAME RATE.

Chemical equilibrium can only be estabilished in a closed system.

A series of reactions in an open system can only come to a steady state.

E.g. Production and destruction of ozone in stratosphere:

• Ozone production: O + O2 --> O3
• Ozone destruction: O3 -hv-> O2 + O
• Ozone destruction: O + O3 --> O2 + O2

None of these comes to equilibrium but left alone will reach a point when ozone is used as fast is it is being produced, so the concentration stays the same = steady state.

For any reversible reaction - many combinations of equilibrium mixtures possible.

Combinations depend on original concs. of substances and conditions.

Position of equilibrium describes one set of equilibrium concentrations for a reaction:

• If most of the reactants become products before reverse reaction sufficiently increases to establish equilibrium = position of equilibrium lies to the right.
• If little reactants have become products when reverse reaction becomes equal in rate to forward reaction = position of equilibrium lies to the left.

Convention is that forward reaction for a given reaction proceeds left to right.

Backward reaction proceeds right to left.

Reactants are always on the left of the arrow and products are always on the right.

Position of equilibrium can be altered by a change in concentration of solutions, pressure of gases or temperature.

Le Chatelier's Principle = If a system is at equilibrium and a change is made in any of the conditions, the system respons to counteract that change as much as possible.

NB: A catalyst doesn't change position of equilibrium, only the rate at which equilibrium is established.

Concentration

• Increasing reactants = shifts to the right (decreases reactants) and vice versa
• Increasing products = shifts to the left (decreases products) and vice versa

In production of calcium oxide from calcium carbonate: CaCO3 (s) <=> CaO(s) + CO2(g) - carbon dioxide is removed from the kiln to encourage equilibrium position to the right to increase the yield of calcium oxide.

Pressure

• Pressure increases - equilibrium shifts to the side with fewer gas molecules.
• Pressure decreases - equilibrium shifts to the side with more gas molecules.

Temperature

• Temp increases - position shifts in the direction of the endothermic reaction.
• Temp decreases - position shifts in the direction of the exothermic reaction.

Rates of reaction can be affected by:

- Concentration

- Surface area

- Pressure

- Size of particles

- Temperature

- Catalysts

- Intensity of radiation

NB: Reaction kinetics = study of rates of reaction.

Reactions occur when particles of reactants collide w/ certain minimum kinetic energy:

• At higher concentrations and higher pressures, particles are in closer proximity so are encouraged to collide more frequently.
• At higher temperatures, a higher proportion of colliding particles have enough energy to react - more particles can overcome the activation enthalpy barrier.
• Smaller particles of reactant = larger surface area on which reactions can take place so there is greater chance of successful collisions.
• Heterogeneous catalysts = provide surface where reacting particles may make and break bonds.

All these increase the rate of reaction.

NB: Particles can be IONS, ATOMS OR MOLECULES.

Activation enthalpy = minimum kinetic energy needed by a pair of colliding particles for a reaction to occur. 

An enthalpy profile = shows how the enthalpy changes as reaction proceeds.

Rates of reaction depend on frequency of collision and energy the particles have when they collide.

Particles need a certain amount of kinetic energy = activation enthalpy.

Activation energy barrier = energy needed to overcome the energy barrier.

As temp increases, the rate of chemical reactions also increases.

Maxwell-Boltzmann distribution = distribution of energies among reacting particles (changes as temp increases).

There needs to be enough molecules with sufficient energy for a reaction to take place - molecules need to have a combined energy greater than the activation enthalpy.

Reactions = faster at higher temps - larger proportion of colliding particles have minimum energy to react.

For any chemical reaction to proceed, bonds need to be broken before new bonds can be made.

Bond breaking = endothermic, bond making = exothermic.

A pair of reacting molecules need to collide w/ combined energy greater than activation enthalpy for a successful collision to occur. Catalysts = used to overcome energy barrier more easily by lowering activation energy barrier.

Catalysts provide an alternative pathway for breaking and making bonds that has a lower activation enthalpy than the uncatalysed pathway.

Catalysts don't affect the position of equilibrium in a reversible reaction.

• Heterogeneous catalysts = provide surface on which reaction can take place - lower energy for successful collision, are in a DIFFERENT STATE to the reactants.
• Homogeneous catalysts = form an intermediate compound w/ reactants, a diagram showing this has two humps, one where the intermediate is formed and one where it is broken down and the product is formed - are in the SAME STATE as the reactants.