The Atmosphere

The Atmosphere

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  • Created by: R_Hall
  • Created on: 28-04-13 14:55

5.2 Molecules and Networks

  • Carbon and silicon are both in Group 4, but CO2 and SiO2 are very different molecules, because of their different bonding
  • CO2 is a gas, because the bonds between molecules are very weak, so little energy is needed to separate molecules. It is soluble in water, giving an acidic solution
  • SiO2 is an extended network of SiO4, where a central Si atom is bonded to 4O atoms. Because of it's extended network structure (giant structure), it is insoluble in water and has high melting/boiling point. The strong covalent bonding means that lots of energy is needed to break the bonds
  • Two types of covalent structure- covalent molecular structure (discrete molecules eg. CO2) and covalent network structure (giant repeating lattices of covalently bonded atoms eg. SiO2)
  • Diamond and graphite are giant lattices of carbon atoms. In diamond, a single C atom is joined tetrahedrally to 4 other C atoms. Symmetrical structure and strong bonds makes it hard. 
  • Graphite is made up of flat layers of C atoms. There are strong covalent bonds between the C atoms in each layer, with delocalised electrons between the layers. This allows graphite to conduct electricity
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6.2 What Happens When Radiation Interacts With Mat

  • When chemicals meet electromagnetic radiation, they absorb energy and this causes the chemical to change. Depends on chemical and energy involved. 
  • Molecule has energy associated with several aspects of behaviour- translation (moving as a whole), rotation (rotating as a whole), vibration (of bonds), electrons. Increasing energy
  • The electronic energy of an atom/molecule changes when an electron moves from one level to another- energy is quantised (fixed levels). All other types of energy are quantised too
  • Molecules change in vibrational energy states when they absorb infrared radiation. The spacing between vibrational energy levels corresponds to the infrared part of the spectrum
  • Making molecules rotates requires less energy- corresponds to microwave radiation (lower frequency)
  • Spacings between translational energy even smaller- can be seen as continuous
  • Electronic changes in a molecules require more energy. Involves jumping electrons to higher enegy levels, corresponding to visible and ultraviolet
  • When a chlorine molecule absorbs 1. electrons excited to higher energy level (later fall back0 2. If higher energy radiation, the atoms break apart- photodissociation- creating radicals. 3. v high energy, electrons leave- ionisation
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6.3 Radiation and Radicals

  • Breaking bonds- bond fission. Either heterolytic or homolytic.
  • Heterolytic- both shared electrons go to one atom- becomes negatively charged, and other is positive. Common when a bond is already polar
  • Homolytic- one or two electrons to each atom. Atoms have no overall charge. The unpaired electron makes the atom highly reactive- a radical. Common when electrons equally shared, but polar bonds can (in gas phase or in presence of light)
  • Radicals with two unpaired electrons are biradicals
  • Radicals are reactive as they tend to fill their outer shell by taking electrons from another atom/molecule. Often steal an electron to become stable, but this causes another radical to be made- radical chain reaction
  • Initiation- radicals formed (through photodissociation)
  • Propagation- reactions destroy and creat new radicals- propagate the reaction
  • Termination- radicals taken out of circulation
  • Radical chain reactions-
  • 1. Often occur in gas phase or in non-polar solvent
  • 2. Initiated by heating or light
  • 3. Go very fast
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7.1 Chemical Equilibrium

  • Dynamic equilibrium occurs when the rate of forward and backward reaction are balanced, in a closed system. It is a reversible change, and concentrations of reactants and products are constant
  • The term position of equilibrium describes one particular set of equilibrium concentrations for a reaction. If the concentration of one changes, the system is no longer in equilibrium and new position of equilibrium will have to be found
  • Position of equilibrium altered by changing - concentration of reacting substances (if solution) - pressure of reacting gases - temperature
  • Le Chatelier's principle- If a system is at equilibrium and a change is made in any of the conditions, then the system responds to counteract the change as much as possible
  • In the case of a change to concentration/pressure, the system at equilibrium moves to counteract the change as much as possible. Eg. if concentration of reactants is increased, the equilibrium moves to product side
  • Heating moves equilibrium to endothermic reaction side, cooling moves to exothermic side
  • A chemical equilibrium can only happen in a closed system, sealed off from surroundings. In open systems, a steady state can be reached- where concentrations are constant, but not an equilibrium (reactants used up as soon as they arrive)
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10.1 Factors Affecting Reaction Rates

  • Reaction rate affected by-
  • 1. Concentration and pressure- when there are more particles, they are closer together, and have a greater chance of colliding with each other, and reacting. 
  • 2. Temperature- at higher temperatures, a much larger proportion of colliding pairs have enough energy to react
  • 3. Particle size- smaller particles have a larger surface area exposed for reaction to take place one
  • 4.Presence of catalyst
  • 5. Intensity of radiation- if involves radiation
  • Collision theory- reactions occur when particles collide, provided that they collide with sufficient kinetic energy
  • A graph of reaction progress against enthalpy is known as an enthalpy profile. Only pairs of atoms with enough KE on collision to overcome the energy barrier, or activation enthalpy, for the reaction will go on to produce products
  • Activation enthalpy- the minimum KE required by a pair of colliding particles before the reaction must occur. Must be supplied to enable bonds in reactant to stretch and break as new bonds form in product
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10.2 The Effect of Temperature on Rate

  • When the temperature is raised, particles move faster, so they collide more frequently
  • Collision theory says that reactions occur when molecules collide with a certain minimum kinetic energy. The more frequent the collisions, the faster the reaction. This energy is needed to overcome the activation enthalpy (energy barrier)
  • As temperature increases, more molecules move at higher speeds, and have higher kinetic energies. 
  • Reactions go faster at higher temperatures because a larger proportion of the colliding molecules have the minimum activation enthalpy needed to react
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10.6 How do Catalysts Work?

  • Catalysts speed up chemical reactions by providing an alternative reaction pathway for the breaking and remaking of bonds that has a lower activation enthalpy. Now the barrier is lower, more pairs of molecules can react when they collide- faster reaction
  • Catalysts do not affect the position of equilibrium- they alter the rate at which it is attained
  • Homogeneous catalysts work by forming an intermediate compound with the reactants
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CS What is in the air?

  • Gases present in the atmosphere- N2 78%, O2 21%, Ar 1%, CO2 ).0383% or 383 ppm (parts per million)
  • To convert from ppm to %, divide the ppm value by 10,000
  • Pollutants and their sources-
  • CO2- combustion of hydrocarbon fuels (eg. in power stations, motor vehicles), deforestation
  • CH4- cattle farming, landfill sites, pice paddy fields, natural gas leakage
  • N2O- fertilised soils, changes in land usage
  • CO- incomplete combustion of hydrocarbons
  • NOx- combustion of hydrocarbon fuels (from reaction of N2 and O2)
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CS Ozone: A vital sunscreen

  • Ozone is present in the atmosphere in tiny amounts. In the stratosphere, it protects us from absorbing dangerous uv.
  • Formed when an oxygen atom (radical) reacts with a diatomic oxygen molecule. Energy to break an O2 can be provided by uv. Some ozone is made in photochemcial smogs- O atoms produced by action of sunlight on NO2
  • The reaction of O2 and O produces O3. The absorption of uv in the splitting of O3 is responsible for the screening effect of O3
  • If rate of producing ozone= rate of destroying ozone, the concentration of ozone would be constant- steady state
  • Scientists predicted the conc of ozone, but it was much lower than the measured value. Suggests that ozone is being removed faster than expected
  • Chlorine and bromine atoms (produced when solar radiation splits CH3Cl and CH3Br in stratosphere). They react with the O3, destroying it. They react much faster with O3 than O atoms do, so have an extremely damaging effect
  • Chlorine atoms act as a catalyst for the reactions, a single Cl can remove about 1 million O3, even small concentrations have a harmful effect
  • Ozone is also removed by hydroxyl radicals (HO) and nitrogen monoxide (NO)
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CS The CFC Story

  • CFC's were first proposed as a refrigerant in 1930. Inventor Thomas Midgley inhaled a lungful of CCl2F2 and used it to blow out a candle, demonstrating it's lack of toxicity and flammability. It replaced ammonia (toxic and smelly) as a widely available refrigerant
  • In the 1970's, two scientists wanted to find out how quickly they would break down in the atmosphere; CFC's were thought to be very stable
  • They looked at all the processes that could affect CFC's in the trophosphere, and calculated how rapidly they could have happen. The answer was very slowly, but they knew that eventually they would break down with fierce UV, producing Cl atoms
  • They found that Cl atoms would most likely react with O3, producing ClO, which would react with O to produce Cl, and the cycle would start again
  • They then calculated how much ozone would be lost, and the resulting figure was much higher than expected. Scientists working in Antarctica were the first to discover the hole in the ozone layer. NASA satellites were using computers to monitor ozone levels and other things, but the concentration of O3 was so low that it was dismissed as an outlier
  • Only in 1990 were CFC's fully banned. The problems were detected so late as the CFC's were so unreactive that the basic technology of the day could not detect the minute concentrations of compounds in the air
  • HFC's have replaced CFC's- no ozone depleting but contribute to greenhouse effect
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CS The Greenhouse Effect

  • The Sun emits energy to the Earth mainly as visible and ultra-violet radiation. The Earth absorbs some of the energy, causing it to heat up and radiate infrared radiation
  • Greenhouse gases absorb some of the infrared radiation emitted from the Earth, and prevents it being re-radiated into space. Methane is a greenhouse gas
  • When methane molecules absorb IR, two things can happen-
  • 1. Some IR is re-emitted in all directions- some back to Earth and some into space
  • 2. Absorption increases vibrational energy of the molecules and causes bonds to vibrate more vigorously. The energy is transferred to other molecules (O2 and N2) by collisions- increasing KE and air temperature
  • CO2 and other gases are greenhouse gases, they let the Sun's visible in, but stop some of the Earth's IR from leaving. Increased concentration of these gases lead to an enhanced greenhouse effect
  • Evidence shows that CO2 and Earth surface temperature have both increased rapidly in the past 100 years, suggesting that more CO2 raises the temperature. Infrared spectroscopy is used in research sites all over the world to determine CO2 concentrations
  • To control CO2 emissions- burn fewer fossil fuels (more C neutral fuels), increasing photosynthesis, burying or reacting in the oceans
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