Module 4 – CHEM Notes PDF

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MODULE 4: Drivers of Reactions MEASURING TEMPERATURE CHANGES  Thermometer/temperature probe Bulb of thermometer immersed in substance to measure temperature  Cup calorimeters Measure temperature changes during a chemical reaction. Styrofoam cup is used to prevent heat being absorbed or released into outside surroundings.  Metal calorimeters Metal has a lower specific heat, and a good insulator  Bomb calorimeters Used when investigating combustion reactions. Takes place in a sealed metal container (the bomb vessel) HEAT, TEMPERATURE, AND SPECIFIC HEAT Heat: Total energy of the moving particles Temperature: a measure of the average energy of that motion   

Some materials absorb heat more readily than others Different volumes of water, heated with identical heat sources for the same amount of time will have different temperature. Different volumes at the same temperature will have different amounts of heat

Specific heat capacity: the amount of heat energy needed to raise the temperature of 1 gram of a substance by one degree Celsius/kelvin. The high heat capacity of water is important to regulate the climate on earth as water is slow to absorb and release heat.

q=mc ∆ T DISSOCIATION OF IONIC COMPOUNDS Dissociation: Refers to the process that occurs when ionic compounds dissolve in water (physical change) Ionic compounds dissociate in water because water is polar molecule. The positive and negative ions of ionic compounds can form ion-dipole bonds with the polar water molecules. MOLAR HEAT OF DISSOLUTION The heat energy absorbed or released when 1 mole of a solid is dissolved in a large quantity of water. Endothermic: Energy used in breaking bonds is more than energy released in making new bonds Exothermic: energy used in breaking bonds is less than energy released in making new bonds. ENTHALPHY Total energy involved in the whole system (H) Change in enthalpy ( ∆ H ): negative for exothermic, positive for endothermic

HEAT OF COMBUSTION Refers to energy released when a substance undergoes complete combustion. Combustion always releases heat, therefore exothermic Thermochemical equation of combustion: includes a balanced chemical equation with the change in enthalpy of the combustion reaction. OTHER REATION ENTHALPIES  Enthalpy of neutralisation: energy change that occurs when acid neutralises a base in aqueous solution  Enthalpy of vaporisation: enthalpy change when 1 mole of a substance changes from liquid to gas  Enthalpy of fusion (melting): heat change when 1 mole of a substance changes from solid to liquid  Enthalpy of atomisation: energy change when one mole of free gaseous atoms is formed from an element in its standard state. C(s)  C(g) ½ O2(g)  O(g)  Enthalpy of formation: energy change when 1 mole of the compound is formed from its elements in their standard state. C(s) + 2H2(g)  C2H4(g) ENERGY PROFILE DIAGRAMS Exothermic: - The products have less energy locked inside them than the reactants - Energy level has dropped ruing the reaction so enthalpy change is negative - Once the activation energy is overcome, the reaction will produce energy and maintain itself. - Ignition temperature is a measure of activation energy. How high the graph reaches up indicates the energy needed to start the reaction.

Endothermic: - The products have more energy than the reactants - Energy level has increased, so the change in enthalpy is positive - Constant energy input is needed to maintain these reactions

THE ROLE OF CATALYSTS  Changes the rate of a chemical reaction without being used up by the reaction  Provides an alternative pathway for the reaction  Type and number of products are not changed but are produced quicker

MECHANISMS FOR CATALYSTS  Forming a chemical intermediate Combine temporarily with one or more reactants and is released to work again and form the product.  Provide a surface on which the reaction occurs The reactants are absorbed temporarily on the surface of the catalyst. This weakens their bonds and lower the activation energy. CATALYST AND ACTIVATION ENERGY Its presence allows more molecular collisions to occur with enough energy to overcome activation energy. CHEMICAL REACTIONS AND BOND ENERGY Bond energy: is a measure of the amount of energy required to break apart one mole of a covalently bonded gas.    

The energy required to break a bond is the same as the energy needed to form that bond A strong bond is a stable bond o Stable bonds form spontaneously, releasing energy in the process The shorter the bond, the higher the bond energy. Bond energy increases as the number of bonds increase

Calculation bond enthalpy: (Energy released when bonds break) – (energy used to break bonds) LAW OF CONSERVATION OF ENERGY The total amount of energy remains constant in any isolated system, which implies that energy can neither be created or destroyed. Exothermic reaction: extra energy has been concerted to heat and perhaps also light and movement. Endothermic reaction: extra energy has been absorbed from the surroundings. HESS’S LAW States that it doesn’t matter whether a product is obtained in one or more steps, the total enthalpy change will be the same using both routes. Two-step pathway When carbon reacts with a limited supply of oxygen it can first form carbon monoxide. ∆ H = -111kJ C(s) + ½ O2(g)  CO(g) If you react carbon monoxide with more oxygen, it will form carbon dioxide. ∆ H = -282kJ CO(g) + ½ O2(g)  CO2(g) Therefore total

∆H

= -393kJ

Applications of Hess’s Law: 1. Using known values of ∆ H to calculate unknown enthalpy values for other reactions This is especially useful for calculating enthalpies of reactions that cannot be determined experimentally.

Enthalpy change for a reaction = the sum of the heats of formation for products – the sum of the heats of formation for reactants

∆ H reaction =∑ ∆ H f ( products )−∑ ∆ H f (reactants)

2. Calculating the energy content of bonds 3. Determining free energy and entropy values 4. Industries can use this information to find out if they are using the most effective process. 5. PHOTOSYNTHESIS AND RESPIRATION Using Hess’s law, we can calculate the enthalpy of the photosynthesis and respiration reaction using enthalpies of formation. Overall endothermic photosynthesis reaction: 6CO2(g) + 6H2O(l)  C6H12O2(aq) + 6O2(g) CO2(g) ∆ H = -393kJ per mole There are 6 moles in the equation so ∆ H = -393 x 6 = -2358kJ H2O(l) ∆ H = -285kJ per mole There are 6 moles in the equation so ∆ H = -285 x 6 = -1710kJ ∑ ∆ H f (reactants) = -2358 + -1710 = -4068kJ C6H12O2(aq) ∆ H = -1275kJ per mole O2(g) is an element, so its heat of reaction = 0 ∑ ∆ H f ( products ) = -1275 + 0 = -1275kJ

∆ H reaction = -1275 – (-4068) = 2801kJ per mole Overall exothermic respiration reaction: C6H12O2(aq) + 6O2(g)  6CO2(g) + 6H2O(l)

∑ ∆ H f ( reactants ) ∑ ∆ H f ( products)

= -1275 + 0 = -1275kJ = -2358 + -1710 = -4068kJ

∆ H reaction = -4068 – (-1275) = -2795kJ per mole SPONTANEOUS REACTIONS Spontaneous process: Takes place without anything being done to make it happen Both exothermic and endothermic reactions can occur spontaneously ENTROPY Entropy (S): the possible ways energy can be distributed in a system of molecules, a measure of disorder. Molecules and energy more spread out/dispersed  increase in entropy Molecules and energy more concentrated  decrease in entropy

Spontaneous chemical reactions happen when their products have more dispersed energy than the reactants. 

Solids have low entropy, especially crystals which have a rigid structure

ENTROPY AND MICROSTATES Microstate: the possible microscopic configurations of a system. The larger the system, and the more particles it contains, the greater the number of possible microstates. A microstate describes the energy of a molecule in a system at any point in time, and the entropy of a system (S) is proportional to the number of possible microstates. MODELLING ENTROPY CHANGES 1. Dispersal model of entropy: We can model the spreading out of matter and energy by placing a crystal of potassium permanganate into the bottom of a large beaker of water and watching as it dissolves and the purple solution spread through the water until it is all uniformly purple

2. Model of moving particles

The molecules rapidly move throughout the whole container at random. The molecules mix spontaneously and no energy change occurs to drive the mixing. The entropy of the system has increased as the molecules and energy have dispersed outwards, becoming less concentrated. 3. Entropy and change of state As a substance changes states from a solid to a liquid, their temperature stays the same but their entropy increases. If the liquid continues to be heated, the same process occurs as they change from a liquid to a gas which further increases entropy. ENTHALPY VERSUS ENTROPY Enthalpy (H) The energy content of a process or closed system. The enthalpy content of the universe is constant. Energy decreases favours a reaction. Measured in joules per kg An expression of first law of thermodynamics Energy is conserved and cannot be created and destroyed.

Entropy (S) The ways energy can be distributed in a system of molecules. The entropy of the universe is not constant – it is always increasing. Entropy (dispersion) increase favours a reaction. Measured in joules per mole per K An expression of second law of thermodynamics Entropy is not conserved.

PREDICTING ENTROPY CHANGES FROM EQUATIONS Entropy increases in the following reactions:  In an exothermic reaction – energy is released increasing temperature of a system or its surroundings.  A reaction that increases the number of moles of gas  The number of moles of solid decreases  Number of particles increase (dissociation)  Pure liquid or solid forms ions  Dissolved gas escapes solution  Larger or more complex molecules Entropy decreases in the following reactions:  Number of moles of gas decreases  Number of moles of solid increases (precipitation)  Number of particles decrease

∆ Suniverse=∆ S system+ ∆ S surroundings ∆ S=∑ ∆ S products + ∑ ∆ S reactants SPONTANEITY AND CHANGES IN ENTHALPY AND ENTROPY Enthalpy tells you about the relative stabilities of reactants and products Entropy tells you how much uncertainty is in the system, how effectively energy is used in the system. Enthalpy (H) Exothermic

Entropy (Ssystem)

Does the reaction occur?

1. Positive (entropy increase)

YES

2. Negative (entropy decrease)

Usually spontaneous, but it depends on T and relative sizes of ∆ H and ∆ S system Usually spontaneous, but it depends on T and relative sizes of ∆ H and ∆ S system

∆ H is negative (energy is released)

Endothermic

3. Positive (entropy increase)

∆ H is positive (energy is absorbed) 4. Negative (entropy decrease)

NO

EFFECT OF TEMPERATURE ON SPONTANEITY If enthalpy is not favourable and entropy is favourable: ∆ G = +H – (T x +S) = +H - TS If T is small ∆ G will be positive, therefore the reaction will not be spontaneous. But if T is big, ∆ G could eventually be negative, which would make reaction spontaneous.

If enthalpy is favourable but entropy is not so favourable: ∆ G = -H – (T x -S) = -H + TS If T is small so TS is small. ∆ G would probably stay negative as TS can’t overcome the -H, so the reaction is spontaneous at low temperatures. So at high temperatures, the reaction is no longer spontaneous. GIBBS FREE ENERGY The amount of energy associated with a chemical reaction that can be used to do work.

∆ G=∆ H −T ∆ S In any spontaneous reaction, Gibbs free energy always decreases – it will have a negative value and be < 0 If you know that the reaction is favoured by both changes in enthalpy ( ∆ H 0 ), then calculating Gibbs free energy is unnecessary. Gibbs free energy calculations is only useful when only one of these conditions – entropy or enthalpy – is favourable, but the other is unfavourable. Gibbs energy is not conserved. EFFECT OF SIGN OF ENTROPY VALUE ON SPONTANEITY If entropy is positive, the reaction becomes more likely to be spontaneous at higher temperatures If entropy is negative, the reaction becomes less likely to be spontaneous at higher temperatures....