Year 12 Chemistry: Unit 4:~ Industrial Chemistry
Chapter 15:~ Fast and Slow Reactions
15.1 Fast and Slow Chemistry
Chemical reactions involve particles (atoms, molecules or ions) colliding and undergoing change during which atoms are rearranged to produce a new particles.
Two examples of different reactions:
J The transfer of electrons: ______
J The transfer of protons (H+): ______
It is important to realise that collisions between reactants so not always result in chemical change. Example: a car engine containing petrol needs a spark (from a spark plug) to initiate the reaction.
READ: Chemistry in Action (pg 247)
15.2 Chemical Energy
What is chemical energy?
The chemical energy of a substance is the sum of its ______energy (stored energy) and ______energy (energy of movement). These energies result from such things as:
Energy changes during chemical reactions
*Draw figure 15.2 a)
Exothermic reactions (-∆H) –release energy
The total chemical energy of the products might be less than the energy of the reactants, therefore of the energy is released into the environment – often as heat energy. Example: combustion of petrol
*Draw figure 15.2 b)
Endothermic reactions (+∆H) – absorb energy
The chemical energy of the products might be greater than the energy of the reactants, therefore energy is absorbed from the environment. Example: photosynthesis.
The energy released or absorbed during a chemical reaction is called the heat of reaction. The heat of reaction is equal to the difference enthalpy between the products and the reactants, it is given the symbol ∆H (delta H).
Enthalpy:______
J Exothermic reactions: H(products) is less than H(reactants), therefore:
-∆H
J Endothermic reactions: H(products) is greater than H(reactants), therefore:
+∆H
Question: 2
Year 12 Chemistry: Chapter 15.2 Chemical energy
15.2 Thermochemical equations
One of the most important endothermic reactions is photosynthesis. Plants absorb energy from the sunlight.
The thermochemical equation for this reaction is:
In this reaction, 2803 kilojoule (kJ) of energy is absorbed when six moles of carbon dioxide reacts completely with six moles of water, producing one mole of glucose (C6H12O6) and six moles of oxygen.
Thermochemical equations show the energy released or absorbed during a chemical reaction.
Energy is measured in joules (J) or kilojoules (kJ).
The heat of reaction, ∆H, has the units J mol-1 or kJ mol-1, to indicate that the amount of energy corresponds to the mole amounts specified by the equation.
The reverse reaction is respiration:
This process results in the release of 2803 kJ of energy for each mole of glucose oxidised.
Another example is the combustion of methanol.
For each mole of methanol fuel that reacts with oxygen 726 kJ of energy is released.
What if the moles are doubled?
If two moles of methanol are used, then twice as much energy would be released and the value of the ∆H is doubled.
Activation energies: getting a reaction started
CH4(g) + 2O2(g) ® CO2(g) + 2H2O(g) ∆H = -890 kJ mol-1
Why doesn’t natural gas burst immediately into flame and release energy when it contacts air? To start a gas oven or heater, must you use a match or electric spark?
What happens to chemical bonds during the course of a reaction:
In the combustion of methane, it results in an overall release of thermal energy. The energy absorbed when the bonds break must be less than the energy released when new bonds form.
The heat of reaction, ∆H, is the net result of the energy absorbed in breaking bonds and the energy released by making them.
The energy required to break the bonds of reactants so that a reaction can proceed is called the activation energy of the reaction,
*Draw figure 15.4 & 15.5
In the combustion of natural gas there is insufficient energy at room temperature to break the bonds of oxygen or methane, a match or spark supplies the energy required to overcome the activation energy ‘barrier’.
An activation energy barrier is present in both exothermic and endothermic reactions.
15.3 Making reactions go faster
The rate at which chemical reactions occur is an important consideration for industrial chemists and chemical engineers. For the manufacture of most products to be profitable, the reactions involved needs to occur rapidly, therefore maximising reaction rates in industry.
Collision Theory
For a chemical reaction to occur, the particles involved must collide with each other with sufficient energy to overcome the activation energy ‘barrier’ to the reaction. This way of visualising reactions is known as collision theory.
It makes sense that the greater the number of collisions between reactant particles, the greater the rate of reaction. In most chemical reactions the majority of collisions do not result in reaction. In these collisions the energy involved is less than that required to break bonds in the reactants, the energy is less than the activation energy. Only ‘successful’ collisions, where the energy of collision is greater than the activation energy, allow a chemical reaction to progress. So, the rate of a chemical reaction is also dependent of the proportion of ‘successful’ collisions.
Factors that affect rates
There are four main ways in which reaction rates can be increased:
Increasing the surface area
In a solid, only those particles that are at the surface can be involved in reaction. Crushing a solid into smaller parts means that more particles are present at the surface. As a consequence of the greater number of exposed particles, the frequency of collisions between these particles and reactant particles increases and so reaction occurs more rapidly.
Example: Sherbet, uses fine grain powders, on contact with the saliva in the mouth, these chemical react quickly.
C4H6O5(aq) + 2NaHCO3(aq) Na2C4H4O5(aq) + 2H2O(l) + 2CO2(g)
(malic acid)
Examples: Fireworks, if larger pieces of aluminium are used instead of fine powder, the reaction is slower and sparks are seen as piece of burning metal are ejected.
Increasing the concentration of reactants
Reactions involving molecules or ions dissolved in solution occur faster if the concentration of the dissolved particles is increased. With more particles moving randomly in a given volume of solution, the frequency of collision is increased and so more successful collisions occur.
Increasing the pressure of gases raises the concentration of gas molecules, causing more frequent collisions. High gas pressures are therefore often employed by engineers in their design of chemical plants that employ gas-phase reactions.
Increasing the temperature
Refrigerators are used to slow down the rate of chemical reactions that lead to food spoilage.
A gardener might use a hot house to accelerate the rate of biochemical reactions in tomato plants to obtain a crop of tomatoes in a short time.
As temperature increases, the average speed and average kinetic energy of the particles increases as well. Hence, as temperature increases so does the rate of reaction.
Extended collision theory
The diagram above shows the distribution or spread of kinetic energies of particles at a particular temperature.
At any given temperature particles have a wide spread of kinetic energies and velocities. Consequently, temperature is a measure of the average kinetic energy.
If we increase the temperature of particles, more particles will move at higher speeds and have higher kinetic energies.
We have seen that a reaction can take place only if the particles colliding have more energy than the activation energy of the reaction. At any instant only a fraction of the particles present have sufficient energy to participate in successful collisions and react.
E∆ represents the activation energy of a reaction. The proportion of particles with kinetic energies in excess of the activation energy is represented by the shaded area used the graph.
As the temperature increases, the shaded are under the graph increases. This means the proportion of particles with more energy E∆ increases. A temperature increase of 10°C almost doubles the number. This factor has a greater impact on rate than the increase in the number of collisions at higher temperatures.
Catalysts
Many reactions occur rapidly in the presence of particular elements or compounds. These substances, known as catalyst, are not consumed during the reactions and therefore do not appear as either reactants or products in reaction equation.
The chemical industry uses catalysts extensively.
Products / Name of process / CatalystPolyethene
Ammonia
Sulfuric acid
Gasoline
Nitric acid
Margarine
Wine
There are two types of catalysts:
J Homogeneous catalyst: are those in the same state as the reactants and products.
J Heterogeneous catalyst: are in different states from the reactants.
How do Catalysts work?
In general, particles at the surface of a solid of high surface energy tend to absorb (from a bond with) gas molecules that strike the surface, lowering the surface energy of the solid. Absorption distorts bonds in the gas molecules or may even break them completely, allowing a reaction to proceed more easily than it would if the solid were absent.
A powdered or sponge-like form of solid catalyst is used to provide the greatest possible surface area. The larger the surface area, the more reactant molecules that can be absorbed, and the faster the reaction.
Haber Process
This industrial process is of great commercial importance as it produces ammonia gas, from which fertiliser nylon, explosives and some pharmaceuticals can be manufactured.
In this process, hydrogen and nitrogen gas are converted to ammonia using a catalyst of powdered iron.
*Draw figure 15.15
N2(g) + 3H2(g) 2NH3(g) ∆H= -91 kJ mol-1
The catalyst allows the reaction to proceed at 500°C and a pressure of 250 atmosphere rather than 3000°C.
The catalyst provides an alternative reaction pathway that dramatically reduces the activation energy ‘barrier’ – the energy needed to break the covalent bonds in the nitrogen and hydrogen molecules.
Consequently a much higher proportion of reactant particles collide with sufficient energy to overcome the activation energy barrier.
So providing such a pathway increases the proportion of successful collisions and therefore the rate of reaction increases.
Note in energy profile that the energy of the reactants and products is unchanged therefore the heat of reaction ∆H, is not changed. Adding a catalyst only changes the activation energy.
Questions: 4, 5, 7, 8, 9, 10, 11, 12, 13, 18 & 20