SCH4U - Unit 5 - Version C - Durham District School · PDF fileSCH4U – Chemistry Lesson...

SCH4U
Grade 12 University Chemistry
Lesson 17 Introducing Equilibrium

SCH4U Chemistry Lesson 17
Copyright 2008, Durham Continuing Education Page 2 of 80
Unit 5: Chemical Systems and Equilibrium Most of the reactions that you have studied in this and other chemistry courses have gone to completion from reactants to products. In many chemical reactions the reaction can proceed in both directions. The steady state of reversible reactions is called equilibrium and it is the focus of this unit. Overall Expectations: demonstrate an understanding of the concept of chemical equilibrium, Le Chateliers
principle, and solution equilibria; investigate the behaviour of different equilibrium systems, and solve problems
involving the law of chemical equilibrium; explain the importance of chemical equilibrium in various systems, including
ecological, biological, and technological systems. This unit consists of four lessons: Lesson 17 Introducing Equilibrium Lesson 18 The Equilibrium Constant Lesson 19 Acid and Bases Equilibrium Lesson 20 Solubility Equilibriums

Lesson 17: Introducing Equilibrium A very important question concerning any reaction is the EXTENT to which it proceeds under some given conditions. Does it proceed to only a limited extent? Or, does it go to completion? Can the reaction be reversed? We can consider the reaction between gaseous hydrogen and chlorine. If hydrogen gas and chlorine gas are heated together, an explosive reaction occurs in which hydrogen combine completely with chlorine to give a quantitative yield (98%) of hydrogen chloride. Thus, this reaction is said to go to completion.
H2(g) + Cl2 (g) 2HCI(g). On the other hand, if gaseous hydrogen and iodine are heated together, the reaction is much less vigorous and does not go to completion. The yield of hydrogen iodide (HI) is not as quantitative as some H2 and I2 remain unreacted.
H2(g) + I2 (g) 2HI(g). We say that the system has attained a state of Equilibrium. Equilibrium will be the focus of this unit.
What You Will Learn After completing this lesson, you will be able to illustrate the concept of dynamic equilibrium with reference to systems such as
liquid-vapour equilibrium, weak electrolytes in solution, and chemical reactions; identify, in qualitative terms, entropy changes associated with chemical and physical
processes; describe the tendency of reactions to achieve minimum energy and maximum
entropy; Recognizing Dynamic Equilibrium
Many chemical reactions can easily run in both forward (towards products) and reverse (towards reactant) directions and are called reversible reactions. This reversible behaviour of the chemical reactions has become known as a reaction "coming to equilibrium."

In a chemical system that can come to equilibrium, both the forward reaction direction and the reverse reaction direction will run all the time. This is the meaning of the word "dynamic" in the title.
The exact moment of equilibrium happens when the rate of the forward reaction equals the rate of the reverse reaction.
When a chemical system is at equilibrium, there are no visible changes in the system. The concentrations of every substance in the reaction will remain constant at equilibrium.
Equilibrium occurs when opposing changes to a closed chemical system occur simultaneously at the same rate. A closed system is a system that may exchange energy but not matter with its surroundings. Conditions that Apply to All Equilibrium Systems There are four conditions that all equilibrium systems: 1. Equilibrium is achieved in reversible process when the rates of opposing changes
are equal. A double ended arrow () indicates reversible changes. For example:
H2O (l) H2O (g) 2. The observable (macroscopic) properties of a system at equilibrium are constant (for
example, colour, pressure, concentration, pH) 3. Equilibrium can be approached from either direction 4. Equilibrium can only be reached in a closed system. For example, carbon dioxide in
a soda drink is in equilibrium when the bottle is closed.
H2O(l) H2O(g)
H2(g) + Cl2(g) 2HCl(g)

Solubility Equilibrium is a dynamic equilibrium between a solute and a solvent in a saturated solution in a closed system. Dissolution is the process of dissolving. Consider orange drink crystals dissolving in water. The orange crystals will dissolve into solution as the solute molecules collide with the water molecules. As the orange crystals dissolve, the solution becomes more and more orange. Even when the solution is saturated, the crystals and dissolved ions are still moving. As the dissolved ions hit the solid drink crystals, they may form ionic bonds and crystallize out of solution. Nearing equilibrium, the rate of dissolution and crystallization approach one another and no visible changes are present. Phase Equilibrium - A dynamic equilibrium called phase equilibrium can exist between different states of a substance in a closed system. In a closed system, molecules in the liquid phase can gain enough energy from collisions to move in the gaseous phase. Similarly as the number of gaseous molecules increases, these molecules can move back into the liquid phase. Chemical Reaction Equilibrium - Some chemical reactions called quantitative reactions proceed completely, meaning all of the reactants are converted into products based on the balanced chemical equation. Consider the decomposition of calcium carbonate in an open system: CaCO3(s) CaO(s) + CO2(g) In this open system, equilibrium cannot be established since the carbon dioxide gas escapes into the external environment as the calcium carbonate decomposes. If the same reaction was placed in a closed container and allowed to proceed, both reactants and products are present after a given time frame:
Blue anhydrous cobalt(II) chloride and pink hydrated
The effect of temperature on the endothermic, aqueous equilibrium: Co(H2O)62+ + 4Cl- CoCl42- + 6 H2O Pink Blue The violet solution in the center is at 25C and contains significant quantities of both pink Co(H2O)62+ and blue CoCl42-. When the solution is cooled, it turns pink because the equilibrium is shifted to the left. Heating the solution favors the blue CoCl42- ions.

CaCO3(s) CaO(s) + CO2(g) (Notice the arrow indicates both the forward and reverse direction) Allowing the reaction to reach equilibrium limits the amount of product produced. Lets consider another reaction in a closed system that can also reach equilibrium N2O4(g) 2NO2(g) Suppose you conducted two experiments, one which began with N2O4(g) molecules in a closed reaction vessel, and another that began with NO2(g), in a reaction vessel. Would the system still reach equilibrium in both cases? Table 17.1: N2O4/NO2 equilibrium concentrations Initial Concentrations (mol/L) Final Concentrations (mol/L)
N2O4(g NO2 N2O4(g NO2 Experiment 1 0.75 0 .721 0.058 Experiment 2 0 1.50 .721 0.058 According to the data above, equal ratios of N2O4(g) and NO2(g) were present regardless of whether the reaction started in the forward or the reverse direction. Also note that equilibrium does not mean that the concentrations are equal, rather that the rates of the forward and reverse reactions are equal. Percent Reaction at Chemical Equilibrium Example 1: Consider the following equation for the formation of hydrogen fluoride from its elements at SATP
H2(g) + F2(g) 2HF(g) If the reaction begins with 1.00 mol/L concentrations of H2(g) and F2(g) and no HF(g), calculate the concentrations of H2(g) and HF(g) at equilibrium if the equilibrium concentration of F2(g) is measured to be 0.24mol/L. Solution 1: Lets start off by stating our givens: [H2(g)]initial = 1.00 mol/L [F2(g)]initial = 1.00 mol/L [HF(g)]initial = 0.00 mol/L [F2(g)]equilibrium = 0.24 mol/L

One convenient method for solving equilibrium problems is by using an I.C.E table. I - Initial concentration C - Change in concentration E - Equilibrium concentration So we will begin by setting up an ICE table for this reaction. H2(g) + F2(g) 2HF(g) Initial concentration (mol/L) 1.00 1.00 0.00 Change in concentration (mol/L) Equilibrium concentration (mol/L) By the time equilibrium has been reached, a certain amount of H2(g)and F2(g) will have formed product (HF). However since an equilibrium is indicated by the arrow (), the reaction will not go to completion (completely to the right). Rather at equilibrium, there will be amounts of each reactant and product available. We will assign the variable x to represent the change in concentrations of reactants and products. Assuming that some of the H2(g) and F2(g) we start with will be converted (or lost in the formation of HF), we will assign them the value of x, and the amount of HF formed will be +2x. Notice how the co-efficients for the balanced equation are included. Lets fill in our ICE table: H2(g) + F2(g) 2HF(g) Initial concentration (mol/L) 1.00 1.00 0.00 Change in concentration (mol/L) -x -x +2x Equilibrium concent

SCH4U - Unit 5 - Version C - Durham District School · PDF fileSCH4U – Chemistry Lesson...

Documents

Transcript of SCH4U - Unit 5 - Version C - Durham District School · PDF fileSCH4U – Chemistry Lesson...