Presentation on theme: "AS Chemistry ATMOSPHERE - INDEX PAGE COVALENT STRUCTURES REACTION RATES ENTHALPY PROFILES ACTIVATION ENTHALPY AND TEMPERATURE CHANGES ON REACTION RATE."— Presentation transcript:
ATMOSPHERE - INDEX PAGE COVALENT STRUCTURES REACTION RATES ENTHALPY PROFILES ACTIVATION ENTHALPY AND TEMPERATURE CHANGES ON REACTION RATE CATALYSTS IN PROVIDING OTHER ROUTES TO LOWER ΔA HOMOGENEOUS CATALYIS AND INTERMEDIATE FORMATION LE CHATELIER’S PRINCIPLE AND IMPACT ON HOMOGENEOUS REACTIONS DYNAMIC EQUILIBRIA AND FACTORS AFFECTING IT GASES IN THE ATMOSPHERE PPM AND % CONCENTRATION OZONE DEPLETION BY HALOGENOALKANES HOMOLYTIC/HETEROLYTIC BOND FISSION RADICALS AND RADICAL CHAIN REACTIONS HALOGENOALKANES, RADIACLS AND PHOTODISSOCIATION CFCS THE ROLE OF OZONE AND PROBLEMS WITH IT EM SPECTRUM WITH REGARD TO SUN ENERGY LEVELS (IR ETC…) ΔE=HV (RADIATION WITH MATTER) GREENHOUSE EFFECT GLOBAL WARMING AND INCREASES GAS CONC. CONTROLLING CO2 EMISSIONS
START -- RADIATION AND MATTER RADIATION AND MATTER CHEMICAL IDEAS 6.2 AND 6.3
RADIATION AND MATTER Electromagnetic radiation may interact with matter, transferring energy to the chemicals involved, these changes depend upon the chemicals involved and the frequency of radiation with which the molecule interacted. When atoms absorb energy, electrons in their outer shells become excited and are promoted from ground energy levels to higher ones; they are in an excited state. We also know that this energy is quantised and can be seen as distinct black lines on a bright background in an absorption spectrum. As molecules, chemicals always have energy about them whether it be movement of the entire molecule or vibration of the bonds; it happens. The electrons in the bonds between atoms also have their own electronic energy and may move between energy levels. The molecule’s energy can be associated with… - energy of translation (the movement of the whole molecule) - energy of rotation (the movement of the molecule) - energy of vibration (vibration of the bonds within the molecule) - energy of electrons (electronic energy) As electrons in an atom or molecule move between different energy levels, the electronic energy of the atom or molecule changes. This energy has fixed values; is quantised. ALL FORMS OF ENERGY ARE QUANTISED, WHETHER THEY BE TRANSLATIONAL, ROTATIONAL OR VIBRATIONAL. INCREASING ENERGY
RADIATION AND MATTER Before, we said that all energy is quantised. If we take HCl as an example; it may occupy certain fixed vibrational energy levels. When the HCl molecule absorbs energy of a certain frequency, it’s vibrational energy level increases correspondingly. The energy between vibrational energy levels for HCl is equal to one photon of IR, meaning that infrared radiation can make HCl molecules and many others, increase their vibrational energy. W hen they absorb IR of a certain frequency, they vibrate more violently and so it happens that HCl absorbing IR of a particular frequency (to make it vibrate in a particular way) can be used to identify it. Changes between vibrational energy levels correspond to the infrared part of the EM spectrum. We feel infrared radiation as heat, as it vibrates the bonds within chemicals on and in our skin. The energy of the radiation is therefore transferred to kinetic energy, which raises the temperat- ure. As we know, making molecules rotate requires less energy than needed to make a molecule’s bonds vibrate; as such it corresponds to a lower frequency region of the EM spectrum (as energy is proportional to frequency (E=hv). The region we are speaking of here is the microwave region; the simple chemistry here is the basis of how microwaves and such cooking techniques work. Changes in translational energy are so small that they are treated as being continuous; translational energy lies on a continuum. EE
RADIATION AND MATTER CHANGE OCCURINGSIZE OF ENERGY CHANGERADIATION ABSORBED ROTATIONAL 1 × 10 ⁻²² > 1 × 10 ⁻²⁰ MICROWAVE VIBRATIONAL 1 × 10 ⁻²⁰ > 1 × 10 ⁻¹⁹ INFRARED ELECTRONIC 1 × 10 ⁻¹⁹ > 1 × 10 ⁻¹⁶ VISIBLE OR UV A table to show the bounds of frequencies absorbed by matter and the effect this has on the molecule. As translational energy is on a continuum, it is assumed to happen without the influence of radiation by measure of the above. Electronic requires the higher frequencies to happen, and so the radiation with the most energy. Other frequencies of radiation cause different changes within an molecule or atom for example; gamma rays cause changes within the nucleus of an atom. Although such cases are not of concern here, in any case the energy absorbed will be detailed as bounds (not finite values) because the amount frequency of radiation required to produce the desired effect will depend upon the chemical and the bonds involved. For instance, C - F bonds are much stronger than C -Br bonds and so require more energy to vibrate, the molecule containing the C - F bond is said to be more photochemcially stable.
RADIATION AND MATTER Just as electrons in an atom occupy definite energy levels and may be promoted by visible or UV light, molecules absorb energy to become promoted to higher levels. Electrons within the molecule, bonding or not will occupy levels which they can be moved into or out of; those in the outer shell may move into higher levels more easily than in lower energy levels. When a molecule absorbs radiation, one of three things may happen. Let’s take chlorine as an example… -Excitation: electrons are promoted to higher energy levels non-permanently and will fall back when the energy is no longer supplied. This is what causes chlorine to have it’s green colour; the Cl2 molecules absorb visible light, except for green. -Dissociation: if enough energy is absorbed, bonding electrons may no longer be able to hold the atoms together and the molecule may separate (dissociate). This will often produce radicals that due to their high tendency to react, will react again to form more radicals in a radical chain reaction (where possible). When is visible light or UV that causes this dissociation, it is called photodissociation. The formation of chlorine radicals by this process is a main driving force in the destruction of the ozone layer. -Ionisation: when enough energy is absorbed, electrons may be promoted to levels where they are no longer attracted to the nucleus. The electron leaves the nucleus and the atom is ionised. INCREASING ENERGY
RADIATION AND MATTER The energy of a single photon of radiation of a particular frequency is calculated using the formula: E=hv where h is the Planck constant [6.63 ×10 ⁻⁶⁴ J/Hz] and v is the frequency of the wave. CHEMICAL IDEAS EXAMPLES A beam of infrared radiation has the energy or 3.65 × 10 ⁻ ² ⁰. Calculate the frequency of the radiation… - E=hv, therefore 3.65 × 10 ⁻ ² ⁰ = 6.63 ×10 ⁻⁶⁴ × v - (3.65 × 10 ⁻ ² ⁰ ) ÷ (6.63 ×10 ⁻⁶⁴ ) = v - v = 5.05279035 × 10¹³ To change one mole of molecular HCl from the lowest vibrational energy level to the next requires 32.7kJ. How much energy in joules is needed to change one molecule of HCl from ground to the next vibrational energy level. What frequency would be absorbed when HCl absorbs energy in this way, what type is it and what it it’s wavelength (use: c=λv)… - 32.7kJ = 32,700j per mole 32700 ÷ 6.02 × 10²³ = 5.431893688.. × 10 ⁻ ² ⁰ - x.y × 10 ⁻ ² ⁰ means infrared the v of this radiation… (5.43… × 10 ⁻ ² ⁰ ) ÷ (6.63 × 10 ⁻ ³ ⁴ ) … v = 8.19… × 10¹³ - c(3.00 × 10 ⁸ )= λv - 3.00 × 10 ⁸ = λ × 8.19… × 10¹³ - λ = (3.00 × 10 ⁸ ) ÷ (8.19… × 10¹³) - λ = 3.66 × 10 ⁻⁶ m
HETEROLYTIC AND HOMOLYTIC BOND FISSION Chemical reactions involve the breaking and reforming of bonds and the way in which this happens has an important influence on reactions. A covalent bond shared between two atoms, in HCl for example, will consist of a pair of electrons shared between these two atoms. When this bond breaks in a reaction, this bond may either break homolytically or heterolytically depending upon the energy supplied and conditions. Heterolytic fission happens when the two bonding electrons go to only one of the atoms, forming an anion and subsequently a cation. In HCl, the hydrogen becomes the cation, whilst the chlorine becomes the anion due to it’s electronegativity. This type of fission is common where a bond is already polar, as the one mentioned above. The halogenoalkane bromomethane has a polar carbo-halide bond which may break heterolytically. Homolytic fission is the fission whereby the paired bonding electrons are distributed evenly between the two constituent atoms, forming (if only for a minute period) two radicals. Due to the high reactivity of this species, both will often react given the opportunity. These reactions happen most commonly when the electrons are evenly distributed across atoms in a bond, that is to say that the bond is more non-polar than not due to similar electronegativities. Any of the C-H bonds in methane may break in this way, forming two radicals, however this type of reaction can occur with polar molecules or such, when reactants are in the gas phase and particularly in the presence of light.
RADICAL CHAIN REACTIONS Radicals tend to be a reactive species because of their unpaired electron, however some radicals exist stable enough to live as molecules for a period of time. Nitrogen monoxide has only seven electrons in it’s outer shell and so is a radical, but it is not as reactive as say, a chlorine radical. Radicals with more than one unpaired electron in their outer shell are called biradicals. The dioxygen molecule (O ₂ ) has two unpaired electrons in it’s outer shell that do not simply pair up because they have opposite spins; they will not occupy the same orbital. Radicals, due to their nature will create radical chain reactions because they will always steal one electron from another bond, leaving behind another unpaired electron that reacts to form another radical and so on. These radical chain reactions follow three stages… - Initiation: the RCR will be initiated by the creation of a radical, this may be caused by photodissociation. - Propagation: the reactions happen en mass, forming more radicals and causing more and more reactions. -Termination: occasionally, two radicals will meet and form a product with no unpaired electrons. This doesn’t happen often or straight away because radicals are not present in the same space for very long at all and will be forming new products in reactions for most of the time. If we exemplify the above process for the production of hydrogen chloride, we can see that the the production of chlorine radicals due to photodissociation initiates our reaction, which is then propagated via H ₂ to produce H radicals and so on… Cl ₂ + hv → Cl◦ + Cl◦ Cl◦ + H ₂ → HCl + H◦ H◦ + Cl ₂ → HCl + Cl◦ [H ₂ + Cl ₂ → 2HCl]
METHANE AND CHLORINE Methane and chlorine react vigorously in the presence of light to produce chloromethane and other products such as ethane and chlorine (which may go on to react further). We know that alkanes are a generally unreactive species with polar solvents or organic solvents, but they will react with halogens in the presence of light. This reaction will not happen in the dark, but as soon as they are exposed to sunlight or light, the radical chain reaction will begin. It is initiated by the photodissociation of chlorine molecules, producing two chlorine radicals. These then go on so react with methane molecules to form hydrogen chloride and methyl radicals which react with chlorine molecules and so on. This propagation is ended when two radicals such as chlorine and chlorine, chlorine and methyl or methyl and methyl meet. The overall reaction is represented in green, and shows the production of chloromethane and hydrogen chloride, however; the presence of chlorine will cause further substitution to form dichloro- and trichloromethane. These are examples of radical substitution reactions. So t o summarise radical chain reactions; they often occur very fast and in either the gas phase or non-polar solvent, and they often require heating or light to start. Cl ₂ + hv → Cl◦ + Cl◦ Cl◦ + CH ₄ → HCl + CH ₃ ◦ CH ₃ ◦ + Cl ₂ → CH ₃ Cl + Cl◦ Cl ₂ + CH ₄ → CH ₃ Cl + HCl CH ₃ ◦ + Cl◦ → CH ₃ Cl CH ₃ ◦ + CH ₃ ◦ → C ₂ H ₆ Cl◦ + Cl◦ → Cl ₂ INITIATION PROPAGATION TERMINATION OVERALL
START -- CHEMICAL EQUILIBRIA CHEMICAL EQUILIBRIA CHEMICAL IDEAS 7.1
CHEMICAL EQUILIBRIUM The term “equilibrium” refers to a state of balance; in chemistry, this definition also encapsulated the dynamics of a state and therefore is called the dynamic equilibrium. In a dynamic equilibrium, the concentration of products and reactants will be at a constant state, being produced and used at the same rate in a closed system. Consider a bottle of carbonated water; the bottle is sealed, creating a closed system whereby neither the liquid nor the gas is escaping. If the bottle is left in a room at a constant temperature for a period of time, the contents will appear to settle and nothing will appear to be happening on a macroscopic scale, however this is not the case. If the pressure of the carbon dioxide gas above the water and the concentration of the gas in solution were to be measured they would read a constant level, but on a microscopic scale (in terms of molecules) there is a lot going on. Molecules of CO ₂ in the air are moving about and colliding all of the time, as are molecules in the water. Molecules that are just above the surface of the water may enter the solution (gas -> aqueous) through the inevitable bumping into of water molecules, and vice versa. The turning of carbon dioxide from gas to aqueous is a reversible change. It is represented by an equals sign consisting of two single headed arrows pointing in opposite directions. Sticking with the idea of the drinks bottle, the water can be said to be in a state of dynamic equilibrium as well as the carbon dioxide. As well as the water being in liquid form in the carbonated drink, it also exists in the gaseous phase above the water, again; the molecules colliding and condensing to become aqueous just as they escape from the liquid. Owing that the cap is still on the bottle, the system is closed and after a while a dynamic equilibrium will be reached.
CHEMICAL EQUILIBRIUM Carbon dioxide, as well as simply changing states is able to react with water molecules on entering the liquid. This only happens to a small proportion of the carbon dioxide molecules that enter the drink, but this small amount go on to react with water to produce hydrogencarbonate ions and hydrogen ions as seen below. It is a reversible reaction. CO ₂ + H ₂ O ↔ HCO ₃⁻ + H ⁺ NO ₂ + CO ↔ CO ₂ + NO Molecules of carbon dioxide in the water are colliding with water molecules all the time to react in reaction above, but if you were to measure the pH of the liquid, it would not change as more hydrogen ions are produced, because whilst carbon dioxide and water are reacting, so are hydrocarbonate ions and hydrogen ions. This is another example of a dynamic equilibrium; gain nothing seems to be changing on a macroscopic level, but microscopically there is much activity. Yet another example finds itself in the reversible reaction between nitrogen dioxide and carbon monoxide. If the two gases are placed in a sealed container together, the reaction above takes place. Initially, there won’t be a reverse reaction because there are products to convert back into reactants but as the concentration of carbon dioxide and nitrogen monoxide increases, so will the rate of the reverse reaction. After a time, a point will be reached when the rate of the forward reaction is equal to that of the reverse; the concentrations of both products and reactants will be equal. Once again we have reached a dynamic equilibrium. EQUILIBRIUM NO2 + CO NO + CO2
POSITION OF EQUILIBRIA There are many possible equilibrium mixtures in a given reaction system, that are affected by changes to the conditions in the system. The position of the equilibrium is used to describe the concentrations for a reaction, if one of the concentrations is changed then the position of the equilibrium will move to compensate for this (the Le Chatelier’s principle). When concentrations are changed, or indeed external factors like temperature, pressure and concentration, then the system is no longer in equilibrium and the concentrations of substances will change until a new equilibrium is reached. When this equilibrium is reached, the rate of forward and backward reactions are the same but the concentrations of products and reactants are unlikely to be. In the case where most of the reactants have been converted into products before the reverse reaction reaches equilibrium, we say that the equilibrium lies to the products. When this is the other way around and little reactants are converted before equilibrium is reached, the equilibrium lies to the reactants. Once a system has reached equilibrium, it is impossible to know whether the reaction was initially started with products or reactants because either way you will reach the same rates of reaction and therefore equilibrium (given the same conditions).
LE CHATELIER’S PRINCIPLE The position of an equilibrium can be changed by the following, note that a catalyst does not change the position of the equilibrium, it merely lowers the activation enthalpy for the forward and reverse reactions, thus making the reaction achieve equilibrium more quickly… - concentration of reacting substances (if in solution) - pressure of reacting gases - temperature If we take yellow Fe(III)³ ⁺ ions and colourless thiocyanate ions (SCN ⁻ ), they react in solution to give us a deep red [Fe(SCN)]² ⁺ in a reversible reaction. The intensity of the red colour gives us an indication of the position of the equilibrium; if the red is intense, then the equilibrium lies towards the products as there is a high concentration of iron thiocyanate ions present. If the colour is weak, then we can assume that the equilibrium lies to the reactants as little iron thiocyanate is being produced. Increasing the concentration of iron or thiocyanate ions intensifies the red colour, showing a shift to the right. If we add ammonium chloride, chloride ions react to form FeCl ₄ making the red paler as the equilibrium shifts to produce more iron ions. From this we can draw the following statements; increasing the concentration of the reactants causes the equilibrium to shift to the right to make more products, decreasing the concentration of reactants causes the equilibrium to shift to the left to make more of them; increasing the concentration of the products causes the equilibrium to shift to the left to make more reactants and if we decrease the concentration of products then the equilibrium will shift to the right to make more of them.
POSITION OF EQUILIBRIUM Changing the pressure of a reaction is important only when we consider reactions that involve gases, homogeneously or heterogeneously. In industry there are many important processes that need the equilibrium to be shifted as far to the right as possible and for these reactions, the exact conditions under which such an affect can be bad must be known. In general we know that; increasing the pressure in a system causes the equilibrium to shift to the side with fewer molecules as to lower the pressure, decreasing the pressure causes the equilibrium to move to the favour the reaction that produces most molecules, to raise the pressure. This is consistent with le Chatelier’s principle. If we take the steam reforming of methane, we can see that an initial decrease in pressure favours the forward reaction (production of carbon monoxide and hydrogen) whilst in the following production of methanol, high pressures favour the forward reaction. Supplying heat to a system always increases the rate of reaction due to collision theory. If the reaction is reversible and the forward reaction is exothermic, then the reverse reaction will be favoured. For example with nitrogen dioxide (brown) and it’s colourless dimer dinitrogen tetraoxide, the forward is exothermic, the reverse; endothermic. In a sealed container, heating causes the mixture to turn browner and cooling has the opposite effect. In general we say that; heating a reaction causes the equilibrium to shift towards whichever reaction is endothermic, taking in the heat and reducing the temperature; cooling the mixture will cause the equilibrium to shift towards the exothermic reaction to create heat and raise the temperature. Again this is consistent with le Chatelier’s principle. WHENEVER A ∆H IS GIVEN FOR A REVERSIBLE REACTION, IT IS ACCEPTED THAT THIS IS FOR THE FORWARD REACTION ONLY.
EQUILIBRIA There are a number of systems, particularly in nature where various equilibria depend on or are linked to each other. The reversible reaction between carbon dioxide gas and carbon dioxide in solution facilitates the production of hydrogen carbonate ions and hydrogen ions in the oceans. Le Chatelier’s principle allows us to make observations and predictions that changes will have on certain natural systems such as the one above. For example, increasing the concentration of gaseous carbon dioxide will cause the equilibrium of the first equation to shit to the right and produce more aqueous CO2, meaning there is a higher concentration in the water, causing the equilibrium of the second equation to shift to the right as well, producing more ions and ultimately making the sea water more acidic. CO ₂ (g) ↔ CO ₂ (aq)CO ₂ + H ₂ O ↔ HCO ₃⁻ + H ⁺ Strictly speaking, equilibrium cannot be obtained in an open system as conditions do not remain constant enough for one to be established. This is the case with the decomposition of calcium carbonate (limestone) into calcium oxide (quicklime) and carbon dioxide; if the system is open then the carbon dioxide will escape and the equilibrium will shift to make more, eventually converting all limestone into quicklime. Although an equilibrium may not be obtained, a steady6 state can6 still be achieved whereby the concentrations of reactants and products remain constant. Another example is the production and destruction of ozone in the atmosphere.
START -- REACTION KINETICS REACTION KINETICS CHEMICAL IDEAS 10.1, 10.2, 10.5 AND 10.6
FACTORS AFFECTING REACTION RATES Different reactions have different rates at which they perform. Some reactions such as the burning of fuel in the cylinder of a car engine, or the precipitation of a silver halide from solution happen fast; the rusting of iron and souring of milk happen slowly. The study of these rates of reaction or reaction kinetics is an important section of chemistry that helps scientists to speed up or slow down chemical processes in industry whilst also helping them to make important predictions about mechanisms of gases in the atmosphere (for example). There are a variety of factors that affect the rate of a reaction, for example; -Concentration: measured in moldm ⁻ ³, if the concentration of the reactants is increased, there are more molecules to be converted into product due to an increase in the number of molecules occupying a given volume. Therefore, the rate of reaction is increased if the concentration of reactants increases; with gases the concentration is proportional to the pressure. -Temperature: in nearly all cases, increasing the temperature increases the rate of reaction due to collision theory. -Intensity of radiation: if the reaction requires radiation to begin, the rate of reaction will increase with the intensity of the radiation. For example; the photodissociation of molecular oxygen to form two oxygen radicals requires a specific frequency of UV radiation to begin, if this radiation is intense, the reaction can happen at a faster rate. -Particle size: if the reaction involves solids, then increasing the surface area available will increase the R.O.R. -Catalyst: increases the rate of reaction through lowering the activation enthalpy by providing an alternative route.
FACTORS AFFECTING REACTION RATES Many of the factors affecting reaction rate given on the previous slide are easily explained through the collision theory of molecules. In order for molecules to react, they must first meet and collide but often this interaction simply isn't enough; reacting molecules must also possess a minimum kinetic energy to overcome the activation enthalpy for the reaction. When the concentration of reactants increases there will be more molecules per unit of volume, meaning that there will be a greater number of collisions and therefore successful collisions will happen more frequently (this also goes for increasing pressure). We can say quite definitively therefore that any factor increases the number of collisions increases the rate of reaction as well. As two reactant molecules approach each other and collide, their kinetic energy turns into potential energy and the energy of the reactants rises. If energy of reactants is plotted against reaction progress, this is called an enthalpy profile. Only those molecules with enough energy to overcome the energy barrier will go on to successfully react and form products, and the energy possessed by the reactants depends greatly upon the temperature of the gas. At higher temperatures, a much larger proportion of the reactants have enough energy to overcome the AE. An example of this is the reaction between methane and oxygen, under standard temp and pressure they do not react, but given heat they react explosively. Enthalpy profiles like the one to the left are useful in demonstrating the energy changes over a reaction, but are not representative of all reactions, as many do not simply take place in one step. However, it does show the point in a reaction where old bonds are beginning to break and new ones are forming (highest point). The curve itself shows the pathway of a pair of colliding molecules. ΔAΔA ΔHΔH REACTANTSPRODUCTS Cl + O ₃ → Cl - - - O - - - O - O → ClO + O ₂
TEMPERATURE AND REACTION KINETICS Increasing the temperature is one simple step upon which many chemical processes in industry are heavily dependent, for example; there’d be no Haber process without it! Through measuring reaction rates at different temperatures, it’s clear to us that for most reactions, the reaction rate is doubled for a rise of just 10°c. Earlier we discussed the idea of collision theory and how increasing the pressure can increase the rate of reaction. If we use the Haber process as an example, it becomes clear why this is the case. The reactants are both gases and increasing the pressure means that the molecules are closer together and so collide more frequently and react more often. It is for this reason that the Haber process uses high pressures. Temperature also has an important effect on reactions, as is increases the kinetic energy of molecules. Through this, more collisions occur because the molecules are moving more and so the reaction rate increases, but roughly how much more frequently are the molecules colliding? We know the answer to this because the speed of molecules is proportional to the square root of the absolute temperature. If we increased the temperature of a reaction from 300K to 310K, the average molecular speed would be expected to rise by √ (310 ÷ 300) = 1.016, which is 1.6%. However, the rate can actually increase by up to 300%. It becomes clear therefore that there is more to increasing the temperature than simply the speed of the molecules…the answer here is the energy with which they collide. Raising the temperature gives molecules more energy with which the activation energy barrier can be overcome, thus; increasing not only the number of collisions, but also the number of successful collisions. N ₂ + 3H ₂ → NH ₃
MAXWELL-BOLTZMANN DISTRIBUTIONS At a given temperature, molecules will be moving at different speeds, some slow, some fast. The distribution of the kinetic energies of a gas at a given temperature (i.e. hydrogen and nitrogen molecules) is given by the Maxwell-Boltzmann distribution as shown below. As the temperature increases, so does the speed of molecules and number of collisions. If we say that the activation enthalpy for a reaction is +50kJ/mol (red line), then through in creasing the temperature, we can see that the proportion of molecules with enough energy to collide successfully and react has increased by quite an amount. IT IS WORTH NOTING THAT THE ENERGY DISTRIBUTION FOR THE KINETIC ENERGY OF MOLECULES IS THE SAME AS THE NUMBER OF COLLISIONS WITH KINETIC ENERGY (E). THE SHAPE REMAINS THE SAME. We can summarise the data given in the MB distributions by saying that the rate of the reaction increases at higher temperatures because a larger proportion of the molecules have enough energy to overcome the activation enthalpy for the reaction. The guide that an increase of 10 degrees doubles the ROR is a rough guide, as reactions that have an AE of more than 50kJ/mol will show a greater increase, this is another rule; increasing the temperature has a greater effect reactions with greater AEs. NUM MOLEC. WITH KINETIC ENERGY (E) E NUM MOLEC. WITH KINETIC ENERGY (E) E 300K 310K
CATALYSTS A catalyst is a substance that speeds up the rate of a reaction but can be recovered at the end of this reaction, chemically unchanged from when the reaction began. However, the catalyst may become physically changed during the reaction (i.e. crumble slightly if solid). A reaction whereby the reaction rate is sped up is called catalysis. Catalysts may either be in the same physical state as the reactants (homogeneous catalysis) or in a different state (heterogeneous catalysis). Homogenous catalysis is common where an enzyme is the catalyst, whereas heterogeneous catalysis is important in many industrial processes. When a solid is used to catalyse a reaction of gases or liquids, the catalysis occurs on the solid of the catalyst, therefore a larger surface area is advantageous as the catalysis can happen more quickly and be more effective. The surface area can be increased through finely dividing the solid into a powder, or by using a fine mesh and dispersing the catalyst on that. Zeolites (aluminosilicate materials) are used widely in heterogeneous catalysis in industrial processes such as the cracking of petroleum fractions. Heterogeneous catalysis occurs in five steps, first; the reactant molecules near the catalyst surface and bonds within the molecules begin to weaken and break as new bonds form with the catalyst. Second; bonds within the molecules break completely and bonds are formed with the catalyst (they are adsorbed). Thirdly; new bonds begin to form between reactant particles on the catalyst surface, weakening the bonds between catalyst and reactants. Fourthly; the products are formed as bonds between product molecules form, breaking the bonds between reactant and catalyst. Fifthly; the products diffuse away from the catalyst down the concentration gradient.
HETEROGENEOUS CATALYSIS It is possible for catalysts to become poisoned and inactive when catalyst poisons, which are adsorbed more strongly to the catalyst surface, prevent reactant molecules from adsorbing to the surface. In industry this is a big problem and feedstocks must be refined before use to prevent poisoning. For many reactions once the catalyst has become poisoned and inactive, is cannot be reused and must be replaced. Many toxins work in this way; inhibiting enzyme-catalysed reactions. Sometimes it is possible to regenerate or clean a catalyst, for example; zeolites used in the cracking of long-chain hydrocarbons can become coated in soot as carbon is deposited on the surface of the catalyst. These zeolites can be continuously recycled in a separate container where hot air is blown at the catalyst, reacting with the carbon to produce carbon dioxide. C = C H H H H H H H H H - H C = C H H H H H - H HHC - C H H H H C = C H H H H H - H
CATALYSTS AND ALTERNATIVE PATHWAYS In order for a reaction to take place, bonds within reactant molecules must first stretch and break so that new products molecules can form. The breaking of bonds is an endothermic process; reactant molecules must first have enough energy to collide and then energy to overcome the activation energy for the reaction. If this barrier is high, then very few molecules will have enough energy to react and the reaction will proceed very slowly, if it is too high the reaction will not happen at all. CATALYSTS WORK BY PROVIDING AN ALTERNATIVE REACTION PATHWAY FOR THE BOND BREAKING AND FORMING PROCESS THAT HAS A LOWER ACTIVATION ENTHALPY. With a lower activation energy, a higher proportion of the molecules have enough energy to overcome the AE and so the reaction proceeds faster. However, a catalyst does not affect the position of equilibrium or composition of the equilibrium mixture, only the rate in which it is reached, nor do catalysts increase the amount of product formed or affect the enthalpy change of the reaction.
HOMOGENEOUS CATALYSIS Homogeneous catalysts usually work through the formation of an intermediate, which then reacts to form a product and the catalyst again. This Is why enthalpy profiles for homogeneously-catalysed reactions have two humps; one for each of the reactions. A popular and common example of homogeneous catalysis is in the destruction of ozone in the stratosphere. Chlorine molecules from CFCs and halogenoalkanes undergo photodissociation to form chlorine radicals which then react with ozone. The product of this reaction is chlorine monoxide, which goes on to react with an oxygen radical to then form dioxygen and a chlorine radical; our original catalyst. In this catalytic cycle, chlorine may catalyse thousands of reactions. Most industrial processes involve heterogeneous catalysis as opposed to homogeneous catalysis, despite the fact that the latter of the two is often more specific and controllable. Homo catalysis can produce higher yields of products as less by- products are formed. For example; soluble rhodium is used as a catalyst in the conversion of methanol to ethanoic acid, a reaction that has a 99% percentage yield. ENTHALPY PROGRESS OF REACTION ΔAΔA ΔHΔH
START -- COVALENT STRUCTURE COVALENT STRUCTURE CHEMICAL IDEAS 5.2
CARBON AND SILICON OXIDES Carbon dioxide is the gas that we breathe out. It is essential to life on Earth, as it enables the existence of photosynthetic organisms, and also acts as insulating layer around the Earth, preventing the temperature from plummeting. Silicon dioxide is also essential to life on Earth as it is present in the Earth’s crust and supports growing plants in soil. As both are in group 4 of the periodic table, you might expect that they are similar in terms of properties, but they couldn’t be more different. Carbon dioxide turns straight from a solid (dry ice; turns water acidic) into a gas at -72°c in sublimation, whereas silicon dioxide is a very hard solid with a very high melting point. The dramatic differences between physical properties comes from the differences in bonding between and within atoms. Both form covalent molecules with oxygen, but carbon dioxide consists of only molecules of carbon double bonded to two oxygen atoms, the intermolecular forces are only weak and therefore it turns easily from a solid to a gas (little energy input is required to break the intermolecular forces. The reason for the double bonding and therefore overall properties of carbon dioxide is the fact that carbon atoms are much smaller than silicon atoms, which are far too large to form double bonds with oxygen. However as a gas carbon dioxide exists as the molecules that it was in as a solid, this is not the same for silicon dioxide. Silicon bonds covalently to four oxygen molecules and so the name “silicon dioxide” is misleading; we should call it by it’s systematic name silicon (IV) oxide. Also called quartz, it is made up of a repeating structure of joined silicon tetraoxide units that each have a half share in the oxygen atoms. Sometimes called a network structure (or giant covalent network), it has a very high melting and boiling point due to the vast amounts of energy required to break the bonds within the molecules. Here, intermolecular bonds are very, very weak but the strength of the covalent bonds and the fact that it has a network structure means that the mp and bp are strong.
COVALENT STRUCTURES So, covalent molecules can exist as either simple covalent molecules or giant networks. Simple molecules consist of only discrete covalently-bonded molecules that will have some form of intermolecular bond; this will determine the state of the compound. For example; carbon dioxide has very weak intermolecular forces of attraction and therefore sublime s easily to a gas. If these intermolecular forces are stronger, the compound is more likely to be a solid or liquid at s.t.p. Covalent molecular substances often dissolve in organic solvents and some, like CO2 dissolve in water. Giant networks have very strong covalent bonds and weak intermolecular forces of attraction, giving them an overall high melting and boiling point. They are insoluble and consist of large numbers or repeating lattices. Though we have only spoken of covalent compounds, covalent elements also commonly exist, for example; hydrogen, oxygen, nitrogen and (more or less) the halogens all exist as diatomic molecules, represented by X ₂ and are also all gases at room temperature. Phosphorous and sulphur exist as bigger covalent molecules; P ₄ and S ₈ respectively. Carbon, silicon and boron may form giant covalent networks which all have high melting and boiling points due to the strong bonds between atoms (carbon may also form simple molecules). Silicon forms large tetrahedral structures in which each silicon atom is bonded covalently to four other silicon atoms, thus creating a tetrahedral bond angle situation.
CARBON; EXAMPLES OF COVALENCY Carbon has the interesting property of being able to form both simple covalent and giant covalent networks. Diamond is made up of carbon atoms bonded covalently to four other carbon atoms to produce a structure very similar to that of silicon (as both are group 4 elements). These C-C bonds are very strong and the tetrahedral shape can extend indefinitely, to create the hardest known substance in nature. Diamond has a very high boiling and melting point, is chemically innate and has symmetry in it’s shape. Graphite is another story however, each carbon atom is bonded only to three other carbons, creating two dimensional layers of carbon with the fourth atom of each carbon being delocalised and moving around in clouds between the layers. The layers are formed of rings of six carbon atoms that are strongly bonded together and due to the delocalised electrons, have slight intermolecular forces of attraction. The delocalised layer means that graphite is capable of carrying a current, which makes it fairly unique for a covalent structure as electrons are usually all occupied. Graphite is a soft, brittle solid whose layers can slip over each other, allowing it to be used as a lubricant and in pencils. Fullerenes were discovered in 1985 when a laser was short at graphite in an inert atmosphere, years later they were made by passing al electric current between two graphite rods. They consist of large numbers of carbon atoms covalently bonded to only themselves, for example C ₆₀ and C ₇₀. Recently, long carbon fullerene tubes have been discovered which have the potential to create super-strong lightweight fibres and filtration devices. Fullerenes are such a recent discovery that their properties are not yet known, we know they can be molecular as C ₆₀ dissolved in benzene to turn it red. Wee know also that C ₆₀ (buckminsterfullerene) is an optical limiter; shining light on it causes it to turn darker instantly.
OZONE PRODUCTION AND DESTRUCTION Our atmosphere is made up of many different gases at different concentrations, which perform different roles; one such gas is ozone. It has the molecular formula O ₃ and a bond angle of 117° and is formed through the photodissociation of dioxygen molecules to form oxygen radicals. These oxygen atoms then go on to react with oxygen molecules to produce ozone, and vice versa. Ozone is present in both the troposphere and stratosphere and is produced AND destroyed in both environments. Ozone first began to form when the first photosynthetic life began to form, producing oxygen that travelled up into the stratosphere and formed ozone. After time, the concentration oxygen increased and the rate of destruction and production of ozone reached a steady state. The general reactions for the production of ozone are shown to the right, though it is also produced through many complex reactions in the troposphere involving photochemical smogs and the dissociation of nitrogen oxides. The reactions to the left show the important UV shielding ability of ozone. The term “steady state” system here is used to describe how the rate of ozone production is roughly equal to the rate of it’s destruction. The destruction of ozone comes about when halide radicals (mostly Cl and Br [which is 100x more effective than chlorine]) that come from the photodissociation of CFCs in the stratosphere react with ozone as above. Other radicals such as the hydroxyl radical and nitrogen monoxide also beak ozone down similarly, and in ALL cases the radical acts as a catalyst in the catalytic cycle, being constantly regenerated in a radical chain reaction mechanism. O ₂ + hv → O◦ + O◦ 0◦ + O ₂ → O ₃ PRODUCTION TERMINATION 0◦ + 0◦ → O ₂ DESTRUCTION 0◦ + O ₃ → O ₂ Cl0◦ + O → O ₂ + Cl◦ Cl◦ + O ₃ → ClO◦ + O ₂
ATMOSPHERIC COMPONENTS OF OZONE DEPLETION Hydroxyl, halide, nitrogenous and other such radicals all contribute towards the depletion of the ozone layer. Hydroxyl radicals exist most commonly as a result of the reaction of oxygen atoms and water molecules. In their radical chain reaction, they react directly with ozone to form hydrogen dioxide and dioxygen, of which hydrogen dioxide reacts with an oxygen atom to regenerate the hydroxyl catalyst and produce dioxygen once more. Nitrogen monoxide will react with ozone molecules to form dioxygen and nitrogen dioxide, which is also a radical! This will then go one to react with oxygen atoms to produce nitrogen monoxide and dioxygen, completing the catalytic cycle. This and the previous hydroxyl reaction therefore not only tale ozone out of the atmosphere, but the oxygen radicals that are required in it’s production. Nitrogen monoxide is found naturally in the atmosphere, but is being produced increasingly frequently in the reaction between dinitrogen oxide and oxygen atoms. Dinitrogen oxide itself is also produced naturally in the breakdown of nitrogen compounds in the soil and from oceanic bacteria, but is also given off from the increased use of nitrogenous fertilisers. The above nitrogenous radicals are different in the fact that they are relatively stable radicals and may be prepared and tested with in a lab, there are actually a number of radicals that behave this way. The others suggested here are more typical radicals, but there are actually hundreds, maybe thousands that contribute to ozone depletion, with many having existed long before humans arrived. However, OUR actions are increasing their concentration and changing equilibria positioning. All above radical reactions are competing with each other, the ones which require the least activation enthalpy will happen more often, those that have a high AE will not.
CFCS [CHLOROFLUOROCARBONS] Concerns over the concentration of certain molecules in the atmosphere and their potential to deplete ozone was first raised in 1970 when nitrogen oxides such as the ones discussed on the previous slide were known to be released in aircraft fumes. This was later found to be inconsequential in terms of jet engine exhausts, as they were not a common transport form, however in 1974 CFCs were recognised and predicted to deplete ozone by Sherry Rowland. It was known that CFCs are very unreactive and stable compounds due to the presence of a halogen in the molecule, and Professor Rowland knew this. On tests to see how they exist in the troposphere, he discovered that their stability existed even there; they reacted very, very slowly and COULD exist unchanged for many years. In his testing he also knew that once they reached the stratosphere, CFCs would be broken down by the fiercely intense UV radiation like most molecules, to produce halogen radicals. On investigating the activity of chlorine atoms, they discovered that chlorine atoms are about 1000 times more likely to react with ozone than anything else, producing chlorine oxide and then more chlorine radicals in a radical chain reaction which is ALSO in this case a catalytic cycle. Professor Rowland tested retested to make sure his calculations were correct, but kept getting the same results; CFCs were responsible for the loss of tonnes of ozone in the stratosphere! To be sure that he was correct before publishing his work in a scientific journal, he cooperated with Hal Johnson to find out that this very chlorine reaction has just been discovered, but without the source of chlorine. Sherry had found it; CFCs. Joe Farman’s work in the Antarctic revealed that a hole in the ozone layer exists, which is exacerbated in the Arctic spring. They discovered this when their initial readings were much lower than expected, and confirmed it when they compiled data with NASA, who had observed the depletion but whose satellites had discarded the data as anomalous, it was that low!
CFCS [CHLOROFLUOROCARBONS] Before CFCs, ammonia was used as a refrigerant due to it’s conveniently low boiling point of -33° and ability to be liquefied and compressed easily. However, it was toxic, stank and often caused problems with leakage. Thomas Midgely was assigned the task of developing an alternative that performed well in compression and evaporation, has a low boiling point so that it could be liquefied, did not freeze easily, was non-toxic, chemically stable and non-flammable. CCl ₂ F ₂ was an ideal solution, belonging to a family called the CFCs that had different boiling points to suit different applications; as refrigerants in air conditioning units, propellants in aerosol cans, blowing agents for expanded plastics and polymers like polyurethane and also as solvents in dry cleaning. CFCs from fridges eventually escaped into the atmosphere when they were scrapped. The trouble with CFCs is their useful properties, most of all their stability. Most chemicals and nearly all organic chemicals are broken down in the troposphere by radicals such as the hydroxyl radical, that act as scavengers. However, CFCs pass straight through this layer and up into the stratosphere where they dissociate and tear up ozone molecules. At the time of their discovery as an industrial chemical though, atmospheric chemistry barely existed, the equipment was not available to monitor the small concentrations that would have been initially present and also; the stratosphere seemed too far away to matter. It is only as technology has advanced that we have been able to observe the true nature of CFCs in the stratosphere. CFCs found appropriate use in aerosol can at around the time of WW2, where they dispensed insecticides. They are good solvents, have low boiling temperatures and are non-flammable, stable and non-toxic; this is what made them a good propellant in cans.
CHLORINE RESERVOIRS AND THE OZONE HOLE The incredibly high tendency to react of chlorine radicals means that they would easily destroy all ozone in the stratosphere if left to do so. The only reason this doesn’t happen is because other molecules in the atmosphere react with chlorine first; for example methane and nitrogen dioxide. Methane is produced naturally by many organisms such as those methanogenic bacteria in rice paddy fields, as well as from human activity. Most of it is oxidised in the troposphere but a significant amount passes on through into the stratosphere where it may react with chlorine atoms to produce an ethyl group and hydrogen chloride (a chlorine reservoir molecule). We call HCl a CR molecule because of it’s ability to store chlorine so that it doesn’t react with ozone. Another reaction that happens is one between nitrogen dioxide and chlorine oxide molecules to form chlorine nitrate (a CR molecule). Some of these CR molecules are carried down into the troposphere and dissolve in rainwater due to their solubility, others stay in the stratosphere and cause problems. Holes and weak spots in the ozone layer form all over the world, but satellite measurements show that large holes form over the Antarctic and to a lesser extent, the Arctic. In the Antarctic, this is due to the special weather conditions that arise during the Antarctic winter and following spring which we will look at now. In the Antarctic Winter, the lack of sun light for months causes the temperature to plummet below -80°c which in turn causes polar stratospheric clouds to form. These clouds consist of tiny solid particles of ice and nitric acid which provide a surface upon which reactions can take place… Cl◦ + CH ₄ → CH ₃ + HCl NO ₂ + ClO◦ → ClONO ₂
CHLORINE RESERVOIRS AND THE OZONE HOLE In this same period a vortex forms as a result of the circulating wind, sealing the cold polar core in a reaction vessel. In this vessel, chlorine reservoir molecules are adsorbed onto the surface of the ice particles and react to form chlorine, which is released as a gas and nitric acid which stays on the particle. In the Antarctic spring, the sun returns and UV radiation breaks down the Chlorine gas to produce once again, chlorine radicals that break down ozone ferociously. Large holes also develop over the Arctic, but less persistent and cold weather means that a stable polar vortex doesn’t form, and so the depletion is not as great as in the Antarctic. However the depletion here affects densely populated areas such as Canada, US and some of Europe. ClONO ₂ + HCl → Cl ₂ + HNO ₃ Cl ₂ + hv → Cl◦ + Cl◦ CFCCl◦ UV 0◦ O₂O₂ Cl0◦ O₃O₃ ClONO ₂ HCl NO ₂ CH ₄ ClONO ₂ + HNO ₃ Cl ₂ + HNO ₃ Cl ₂ Cl◦ UV
OZONE IN THE TROPO AND STRATOSPHERE Ozone is present in both the troposphere and to a larger extent, in the stratosphere, where it has an altogether different role than in the troposphere. Further from Earth, the high levels of UV radiation emitted from the Sun facilitates the dissociation of oxygen molecules and as such, the production of Ozone. This reaction absorbs the UV radiation that would otherwise reach Earth’s surface and cause damage to life, so the production of Ozone is an essential atmospheric mechanism in the prevention of harm to us and other life forms. The ozone in the stratosphere is called the ozone layer and through protecting us from the Sun’s UV radiation, it reduces the risk of getting skin cancer, cataracts and the death of some photosensitive organisms. However in the troposphere, ozone can be a hindrance rather than a benefit. Closer to Earth, tropospheric ozone is formed similarly to in the stratosphere. Photochemical reactions take place between nitrogen oxides, water, oxygen and hydrocarbons, which are both present in exhaust fumes and power station fumes. However, excessive use of nitrogenous fertilisers also causes oxides of nitrogen to be released into the atmosphere and hydrocarbons such as methane are produced by methanogenic bacteria, cattle and sheep excretions, burning fossil fuels, rice paddy fields, vehicle exhausts and leaking North Sea gas. In concentrations over 60 ppb, ozone causes respiratory problems such as bronchitis, aggravates asthma, causes the breakdown of some polymers, causes shortness of breath, lung inflammation, prevents sensitive plants from storing and producing food, causes deterioration of leaves (toxic to plants), reduces yield of crops and life in forests, which could effect local ecosystems. For these reasons, ozone production in the troposphere is seen as a problem and needs regulation. Ozone as well as being a highly reactive gas, is a greenhouse gas that contributes to the greenhouse affect. It plays an important role in the production of hydroxyl radicals which break down many substances in the atmosphere that would otherwise build up and cause problems with health.
OZONE AND CFCS Global cooperation has been in place for a number of years now, to tackle problems that CFCs and ozone depletion not only cause humans, but other life forms such as the plankton in the oceans and the array of food chains that they are involved in. The Montreal protocol established that restricting the production of not on only CFCs but bromine compounds such as the halons used in fire fighting was necessary. Since then the protocol has expanded to include most halogenoalkanes, and the developed world have all but eliminated their use, except for in special instances such as the use in inhalers. As a short term alternative for the use of CFCs are HCFCs, which due to their hydrogen-carbon bond are broken down in the troposphere by HO radicals, creating water and CCl ₂ F (radical) that reacts further. It is only a short term measure however as some HCFCs do travel up into the stratosphere and become photodissociated to form Cl radicals. HFCs are even better in terms of minimising ozone depletion as they have no ozone depletion potential whatsoever. In both replacements mentioned, it is the H-C bond that makes the molecule less stable and so easily broken down in the troposphere. Sadly, both replacements are very much greenhouse gases and so cannot be permanent replacements. Concerns have also been raised about the activity of HFCs in the troposphere as they can break down to produce HF and trifluoroethanoic acid but their concentrations are too small to be felt a problem.