Chemistry Unit 2 Making and using alcohols Making Ethanol: Hydration of Ethene (C2H4) Ethanol can be manufactured industrially by the catalytic hydration of ethene Conditions: Steam & H3PO4 Catalyst H2C=CH2 + H2O --- (H3PO4 + Steam)- CH3CH2OH Only 5% of ethene is converted to ethanol but the gases are recycled and passed back through achieving a 95% conversion rate! Fermentation Carbohydrates are converted into ethanol and carbon dioxide. Is and anaerobic process and carried out in the absence of oxygen. C6H12O6 ------(yeast, 37C)- 2CH3CH2OH Uses of Alcohols

Alcoholic drinks Ethanol as a fuel Methylated spirits Methanol as a clean burning fuel added to fuels to improve combustion

Properties Of alcohols Alcohols have relatively high melting and boiling points. They can form Hydrogen bonds between –OH groups of neighboring molecules. Alcohols also have a lower volatility compared with similar length alkanes. Volatility is the ease that a liquid turns into a gas. Primary Alcohols -OH Group is bonded to a carbon atom with no or one alkyl groups. Secondary Alcohols -OH group is bonded to a carbon atom bonded to two alkyl groups Tertiary Alcohols -OH group is bonded to a carbon bonded to 3 alkyl groups

Combustion of Alcohols In a plentiful supply of oxygen alcohols burn completely to form water and Carbon dioxide. C2H5OH + 3O2 ------ 2CO2 + 3H2O Oxidation of Alcohols Primary and Secondary alcohols can be oxidized using K2CR2O7/H2SO4. During this reaction K2CR2O7 changes colour from orange to green. Primary Alcohols Distillation Primary alcohol + [O] ---(k2cr2o7/h2so4)-- aldehyde Reflux Primary alcohol + [O] ---(k2cr2o7/h2so4)- Carboxylic acid Secondary Alcohols Secondary alcohol + [O] ----(k2cr2o7/h2so4) Ketone Tertiary Alcohols Cannot be oxidized. Esterification An ester is formed when an alcohol is warmed with a carboxylic acid in the presence of an acid catalyst ( H2SO4). CH3CH2COOH + CH3OH --- CH3CH2COOCH3 + H20 Esters Are used as adhesives and solvents in the chemical industry. They are also used in flavors and fragrances. Dehydration of alcohols Conc. H2S04 Reflux 40mins CH3CH2OH ----(CONC H2SO4/REFLUX)-- CH2=CH2 + H20

Halogenoalkanes Are compounds where a halogen atom has replaced at least one hydrogen atom. Reactivity of Halogenoalkanes Halogenoalkanes contain a polar carbon-halogen bond. The polarity arises from the different electronegativites.

Halogen atoms are more electronegativity Electron pair is attracted more towards the halogen atom than carbon The result is a polar bond

The electronegativity of the halogens decreases down the group, resulting in a decrease in polarity of the carbon-halogen bond. Electron deficient carbon atom attracts nucleophiles. (NH3,OH-,H2O) Hydrolysis of Halogenoalkanes

When Halogenoalkanes react with hot hydroxide ions, a nucleophillic substitution occurs. The product is an alcohol.

Aqueous sodium hydroxide is usually used Reaction carried out under reflux

A nucleophile is an electron pair donor Nucleophillic substitution

Rate of Hydrolysis Rates of hydrolysis can be determined using the following experiment; The Halogenoalkane mixture is heated with aqueous silver nitrate, with ethanol added.

Water mixture behaves as nucleophile Ethanol acts as common solvent

As the hydrolysis takes place, halide ions from CH3CH2Cl + H20 --- CH3CH2OH + H+ + Cl- The AgNO3 reacts with the halide ion forming a precipitate of the silver halide.

The rate can be found by calculating the 1/time taken for a precipitate to form. For a controlled test:

Equal amounts in mol must be used Halogenoalkanes with the same chain length Use a water bath to ensure constant temperature

Halogenoalkane Precipitate Colour of precipitate Chloroalkane AgCl White Bromoalkane AgBr Cream Iodoalkane AgI Yellow Two Factors affect the rate of hydrolysis:

Polarity –the carbon-fluorine bond is the most polar so the C atom should attract the nucleophile most readily and gibe the fastest rate of reaction.

Bond Enthalpy – the carbon-iodine bond is the weakest amongst the Halogenoalkanes , so should be broken most easily and give the fastest reaction.

Bond enthalpies

Bond enthalpy is more important than polarity The rate of reaction increases from fluoroalkanes (slowest) to Iodoalkane

(fastest) as the carbon- halogen bond weakens Halogenoalkanes and the environment Halogen containing polymers Polyvinylchloride: used in plastic window frames PTFE: non-stick coatings CFC’s Initially used as aerosol propellants and coolants. Because of their low reactivity, volatility and non- toxicity But break down forming chlorine radicals and deplete the ozone layer. HCFC’s are an alternative.

Infrared spectroscopy Absorption of infrared radiation causes covalent bonds to vibrate. Every bond vibrates at its own frequency. The amount of vibration depends on:

Bond strength Bond length Mass of each atom involved

Modern breathalyzers use infrared spectroscopy when analyzing ethanol in breath. Mass Spectrometry Mass spectrometry is a powerful tool it can be used for:

Identifying Unknown compounds Determining the abundance of isotopes Gain info about structure and properties of compounds

Uses of it today include

Detecting banned substances Analyzing molecules in space Detecting traces of toxic chemicals in marine life

Mass spectra of elements e.g Mg

Three peaks means 3 isotopes Heights give relative abundances Results can be used to find Ar

Spectra tells us that : 79% is Mg-24 10% is MG-25 11% is Mg-26 Ar = (24x79) + (25x10) + (26x11) / 100 = 24.3

Reaction Mechanisms Alkanes – Radical Substitution Describe with the aid of suitable equations, the mechanism for the reaction between methane and bromine in the presence of ultraviolet radiation. CH4 + Br2 --------- CH3Br + HBR Initiation Br2 -------(UV)------ Br + Br Propagation Br + CH4 ---- CH3 + HBr CH3 + Br2 --- CH3Br + Br Termination Br + Br -- Br2 CH3 + Br -- CH3Br CH3 + CH3 C2H6 Further substitutions can take place, as the Bromoalkane can react with further radicals. Forming a mixture of products, CH3Br, CH2Br2,CHBr3,CBr4 Alkenes – Electrophilic addition

Halogenoalkanes – Nucleophillic Substitution

Enthalpy In some reactions the products of the reactions have more chemical energy than the reactants. In others, the products have less chemical energy than the reactants. Exothermic Reactions

The enthalpy of the products is smaller than the enthalpy of the reactants There is a heat loss from the chemical system ΔH has a negative value because heat has been lost.

Endothermic Reactions

The enthalpy of the products is greater than the reactants There is a heat gain in the chemical system ΔH has a positive sign because heat had been gained.

Exothermic reactions Oxidation of fuels:

A common example is the oxidation of methane to form carbon dioxide and water

The products have less enthalpy than the reactants so the excess energy is released as heat.

CH4 + 2O2 ------ CO2 + 2H2O ΔH= -890KJmol-1 Combustion reactions affect us all. We rely on the combustion of petrol and diesel in cars. Respiration Sugars like glucose are oxidized to carbon dioxide and water. C6H12O6 + 6O2 ------ 6CO2 + 6H20 ΔH= -2801KJmol-1 Endothermic reactions Thermal decomposition of Calcium Carbonate. Heat provides energy that drives many reactions. These reactions are usually endothermic. CaCO3 --- CaO + CO2 ΔH= +178KJmol-1

Enthalpy Profile Diagrams Exothermic reactions

Endothermic reactions

Activation Energy Ea Activation energy is the minimum energy required to start a reaction by the breaking of bonds Exothermic reactions Even though the products have a lower energy than reactants, there still has to be an input of energy to break the first bond to start the reaction.

Endothermic Reactions

Standard Enthalpy Changes Standard Conditions

100 kPa (1 atmosphere) -298K (25c) 1 moldm-3

Standard enthalpy change of reaction Is the enthalpy change that accompanies a reaction in the molar quantities expressed in a chemical equation under standard conditions, all reactants and products in their standard states Standard enthalpy change of combustion Is the enthalpy change that takes place when one mole of a substance reacts completely with oxygen under standard conditions, all reactants being in their standard states. Standard enthalpy change of formation Of a compound is the enthalpy change that takes place when one mole of a compound if formed from its elements in their standard states under standard conditions Determination of enthalpy changes An excess of magnesium is added to 100cm of 2.00molddm3 CuSO4. The temperature increases from 20C to 65C. Find the enthalpy change of reaction for the following equation Mg + CuSO4 --- MgSO4 + Cu Specific heat capacity = 4.18gkj Density of solution 1.00gcm Answer

1. Find the energy change 100cm3 of a solution has a mass of 100g temperature change ΔT = (65-20) = +45C Heat gained by surroundings, Q = mcΔT = 100 x 4.18 x 45 J = +18810J So heat lost by chemical system = -18810J

2. Find the amount in mol that reacted

Amount in mol CuSO4 = 2.00 x 100 = 0.200 mol 1000

3. Scale the quantities for 0.200 mol = -18810 For 1 mol (5x 18810) = -94050 = -94.05

Rates of reaction – Collision theory Rates of chemical reactions vary considerably

Combustion reactions have very high rates Rusting of iron has very slow rate

Rate of reaction = change in concentration ÷ time units moldm3s-1 The rate of a reaction is the change in concentration of a reactant or product in a given time Factors that alter the rate of reaction:

Temperature Pressure Concentration Surface area Adding a catalyst

We explain how theses factors effect rate using collision theory

When two molecules collide the reaction will take places if the molecules have sufficient energy to overcome the activation energy and they collide in the correct orientation

Effect of concentration If the concentrations of reactants are increased then the rate will increase!

More molecules per unit volume Molecules will be closer together so there is greater chance of the

molecules colliding Collisions will be more frequent

More collisions will take place and there will be more collisions with an energy greater than the activation energy, so there will be more successful collisions and the rate will increase Effect of Pressure When the pressure of a gas is increased molecules are pushed closer together

The same number of molecules occupy a smaller volume For a gaseous reaction increasing pressure is the same as increasing

concentration

So more collisions will take place and there will be more collisions with the energy greater than the activation energy, there will be more frequent collisions and the rate will increase .

Catalysts Increase the rate of reaction without being used up in the process A catalyst lowers the activation energy of the reaction by providing an alternative route for the reaction to follow.

There are many benefits to using a catalyst: Less energy is required for the molecules to react. Much of this energy comes from burning fossil fuels. If a process requires less energy then this save energy costs. With less energy required less fossil fuels are burnt and less carbon dioxide is released into the atmosphere. Economic Importance of catalysts They allow different reactions to be used with less waste and a higher atom economy. Enzymes as biological catalysts Generate very specific products and operate effectively at close to room temperatures.

Lower temperature and pressures can be used, save money and energy. Forms pure products, no separation needed less expensive Conventional catalysts are usually poisonous and pose disposal problems

The Boltzmann Distribution curve Is the distribution of energies of molecules at a particular temperature, often shown as a graph. The effect of temperature on reaction rate

As the temperature increases, the rate of reaction increases.

More collisions take place, the molecules have more kinetic energy so are moving faster

A higher proportion of molecules have an energy that is greater than the activation energy – more successful collisions.

Thus, more molecules will over come the activation energy of the reaction. There will be more successful collisions in a certain length of time the rate of reaction will increase. The effect of a catalyst on reaction rate

On collision more molecules will overcome the lower activation energy of the reaction. There will be more successful collisions in a certain length of time and the rate of reaction will increase.

Chemical equilibrium Dynamic equilibrium is the equilibrium that exists in a closed system when the rate of the forward reaction is equal to the rate of the backwards reaction. Le Chateliers Principle states that when a system in dynamic equilibrium is subject to a change, the position of the equilibrium will shift to minimize the change The effect of a catalyst on equilibrium A catalyst does not effect the position of the equilibrium because it speeds up both the forward and backward reactions.

The Greenhouse effect Is the process in which the adsorption and subsequent emission of infrared radiation by atmospheric gases warms the lower atmosphere and the plants surface. Green house gases, Water vapor Carbon dioxide

Volcanic eruptions Respiration of animals Burning or decaying matter

Methane

Emitted form production of coal A product of rotting waste in landfill Released form cattle

Gases absorb infrared radiation, this causes their bonds to vibrate. Eventually the vibration molecule re admits this infrared radiation. The effect of a green house gas is dependent on its:

Concentration Ability to absorb infrared radiation

Chemist role in minimizing climate change as a result of global warming by supplying governments with evidence on global wariming. Carbon Capture and storage

storage in old oil fields storage as carbonates

The Ozone Layer The ozone layer is continuously being reformed and broken down. How ozone if formed O2 + UV ----- 2O O + O2 ------ O3 How ozone is broken down O3 + UV --- O2 + O Overall O2 + O <-- O3 Removal of ozone O3 + 0 --- 2O2 Ozone depletion Initiation CFCL3 --- Cl + CFCl2 Propagation Cl + O3 -- ClO + 02 ClO + O -- Cl + O2 Overall O3 + O -- O2 NO + O3 -- NO2 + O2 NO2 + O -- NO + O2

Catalytic Converter 2NO + 2CO -- N2 + 2CO2 a catalyst provides a surface on which the reaction takes place.

CO and NO molecules diffuse over the catalytic surface Adsorbed onto the surface Hold molecules in correct position where they react Molecules are deadsorbed

Green Chemistry

Using renewable resources Saving money Preventing waste Maximizing atom economy Recycling and biodegradability