7. Equilibrium

From Dublin Scioto

7.1 Dynamic Equilibrium

7.1.1 Outline the characteristics of chemical and physical systems in a state of equilibrium.

o Physical Systems - rate of evaporation equals the rate of condensation; particles gain energy and go from liquid to gas; particles lose energy and go from gas to liquid; concentrations of gas and liquid do not appear to change

o Chemical Systems - rate of forward reaction (reactants to products) equals the rate of the reverse reaction (products to reactants); concentrations of reactants and products become constant (particles continue to collide and react); the concentrations are not equal, they just do not change

 Before reactants are mixed, the concentration of the products is 0

 As reactants are mixed, products begin to form (product concentration increases and reactant concentration decreases)

 As products form, the rate of the reverse reaction increases; at some point, the forward and reverse reaction equal one another

7.2 The position of equilibrium

7.2.1 Deduce the equilibrium constant expression (Kc) from the equation for a homogeneous reaction.

For the equation N₂O_4(g) <-> 2NO_2(g)

K_c = [NO₂]² / [N₂O₄]

You have to have brackets and the number in front always goes outside of the brackets.

Instructions: Take the coefficient of the products away, and put the product in brackets. Raise the product to the power of the coefficient. Do the same thing for any other products (they are to be multiplied). Now do the same thing for the ractants. It should be written with the Product values all devided by the Reactant values. (the order dosent matter)

For this equation

N2 + 3H2 <-> 2NH3

It gives the expression:

Kc = [NH3]^2 / [N2] x [H2]^3

7.2.2 Deduce the extent of a reaction from the magnitude of the equilibrium constant.

When Kc >> 1, the reaction goes almost to completion, therefore the forward reaction is favored (more reactants are produced).

When Kc << 1, the reaction hardly proceeds, therefore the reverse reaction is favored (more products are produced).

7.2.3 Apply Le Chatelier’s principle to predict the qualitative effects of changes of temperature, pressure and concentration on the position of equilibrium and on the value of the equilibrium constant.

Le Chatelier's principle makes it easy to predict qualitative effects because the change in one element changes the equilibrium and shifts it either left or right. If heat is added the equilibrium shifts towards the side were heat is not in the reaction, and if heat is taken away then the equilibrium is shifted to the side with the heat to equal out the reactants to products. Increase in pressure favors the side with the fewest gas particles and then when pressure is taken away the side with more gas particles is effected. Increasing the concentration increases the chances for successful collisions and it moves away from the side with the increase in concentration. It is the other way around when you decease the concentration and the equilibrium moves towards the side with the decreased concentration.

7.2.4 State and explain the effect of a catalyst on an equilibrium reaction.

Adding a catalyst has no visible effect or shift on an equilibrium reaction because adding a catalyst increases the rate of both the forward and reverse reactions so there is no net change. A catalyst simply speeds up the reaction, enabling more particles to react on both sides of the equation. A catalyst lowers the activation energy of a reaction.

7.2.5 Apply the concepts of kinetics and equilibrium to industrial processes.

For some chemical reactions certain conditions will yield more product. This is the goal in industrial processes because it means more money. For example, some systems work best when there is heat added to it because the heat excites the molecules and according to the collision theory increases the amount of collisions occurring due to appropriate activation energy and more activity. With the stress of adding heat on the reactants side (endothermic reaction) the shift is toward the products meaning a higher yield which was the goal for industrial processes.

Suitable examples include the Haber and Contact processes.

[[1]] Page 3

[[2]] Slides 18-20

[[3]] Slides 2-4 & 6-8

Key Terms

Contact Process - The process to increase the yield of sulfuric acid (H2SO4). It is the most produced chemical in the world.

Equation:

2SO2 + 02 <--> 2S03 ΔH = -196 kJ / mol

All reactants are gases. Only 2 atm of pressure is needed for this reaction.

The reaction is exothermic, so the temperature is lowered to a cool 450°C, but not too low, or the particles will not react.

The catalyst vanadium oxide lowers the activation energy and required temperature.

Equilibrium - Equilibrium is the balancing of two processes. This usually refers to dynamic equilibrium, which is where the particles continue to react on both sides of the chemical equation, but results in no net change because of equal rates on both sides. Note that when at equilibrium the reaction is still occurring. The net change is zero.

Equilibrium Constant - The measured value of the proportion of reactant to product at a given temperature;altering the temperature will cause a shift in the proportion and the constant will change

Is the ratio of reactants to products it is always the same no matter the mass the only thing that would cause it to change temperature

Haber Process - Processes to increase the yield of ammonia;80% used for fertilizers "the nitrogen fixation reaction of nitrogen gas and hydrogen gas, over an enriched iron or ruthenium catalyst, which is used to industrially produce ammonia."

The production of NH3

Le Chatlier's Principle - When a system in equilibrium is changed, the position of equilibrium shifts in the opposite direction to reduce the effects of the change; equilibrium constant does not change (unless temperature is changed); causes of the shift can be concentration, volume, temperature or pressure; used by industries to maximize yields of desired products.

Sample equation N₂O_4(g) <-> 2NO_2(g)

Adding Products - Causes a shift towards the reactants to maintain ratio

Adding Reactants - Causes a shift towards the products to maintain ratio

Increasing the pressure - Causes a shift towards the side with fewer gas particles due to the molecules becoming more compacted with the increased pressure. In the example the shift would be to the reactants due to the reactants having one gas molecule for the product's 2.

Decreasing the pressure - Causes a shift towards the side with more gas particles due to the molecules becoming less compacted allowing more room for said gas particles. In the example the shift would be towards the products because the products have more gas molecules than the reactants.

Decrease in Volume - Shift to the side with fewer gas particles due to a decrease in space (similar to an increase in pressure) In the example the shift would be to the reactants due to the reactants having one gas molecule for the product's 2.

Increase in Volume - Shift to the side with more gas particles due to an increase in space (similar to a decrease in pressure) In the example the shift would be towards the products because the products have more gas molecules than the reactants.

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