1 / 19

Unit 11

Unit 11. Equilibrium. Dynamic balance. Reactions go to completion when almost all reactants convert to products. (most reactions don’t go to completion) Reversible reactions: occur in both forward and reverse directions (double arrow is used ) N 2 + 3H 2  2 NH 3

henrik
Télécharger la présentation

Unit 11

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Unit 11 Equilibrium

  2. Dynamic balance • Reactions go to completion when almost all reactants convert to products. (most reactions don’t go to completion) • Reversible reactions: occur in both forward and reverse directions (double arrow is used ) N2 + 3H2 2 NH3 • Chemical equilibrium: a state in which forward and reverse reaction balance each other because they take place at equal rates. • Concentrations of reactants and products are CONSTANT but they are NOT necessarily EQUAL.

  3. Equilibrium expressions • Law of chemical equilibrium: states that a given temperature, a chemical system might reach a state in which a particular ratio of reactant and product concentrations has a constant value. aA + bB cC + dD • Equilibrium constant, keq • Keq = [C]c [D]d [A]a [B]b keq > 1 : products are favored at equilibrium keq < 1 : reactants are favored at equilibrium

  4. Homogeneous equilibria • All reactants and products are in the same physical state. • Ex. Write the equilibrium expression for the following reaction: H2 (g) + I2 (g)  2HI (g) • Equation has to be balanced. Keq = [HI]2 [H2] [I2]

  5. Heterogeneous equilibria • Reactants and products are present in more than one physical state. • Keq depends only on the concentration of substances in gaseous state. • Ex. I2 (s)  I2 (g) keq = [I2 (g)] [I2(s) ] *concentrations of solids and liquids are constant so they are combined with keq Keq = [I2(g)]

  6. Ex. SO3(g) + H2O(l) H2SO4(l) • Classwork p601-603 #1-4

  7. Le Chatelier’s Principle • The French chemist Henri Le Chatelier (1850-1936) studied how the equilibrium position shifts as a result of changing conditions • How to control equilibria to make reactions more productive. • Le Chatelier’s principle: If stress is applied to a system in equilibrium, the system changes in a way that relieves the stress. • Stress: any kind of change in a system at equilibrium that upsets the equilibrium

  8. Le Chatelier’s Principle • What items did he consider to be stress on the equilibrium? • Concentration • Temperature • Pressure • Concentration– adding more reactant produces more product, and removing the product as it forms will produce more product Each of these will now be discussed in detail

  9. Le Chatelier’s Principle • Temperature – increasing the temperature causes the equilibrium position to shift in the direction that absorbs heat • If heat is one of the products (just like a chemical), it is part of the equilibrium • coolingan exothermic reaction will produce more product, and heating it would shift the reaction to the reactant side of the equilibrium: • C + O2(g)→ CO2(g) + 393.5 kJ

  10. Le Chatelier’s Principle • Pressure– changes in pressure will only effect gaseous equilibria • Increasing the pressure will usually favor the direction that has fewer molecules N2(g) + 3H2(g)↔ 2NH3(g) • For every two molecules of ammonia made, four molecules of reactant are used up – this equilibrium shifts to the right with an increase in pressure • CW handout 17.2

  11. Using Equilibrium Constants Ex. At a temperature of 1405K, hydrogen sulfide which has a foul odor resembling rotten eggs, decomposes to hydrogen and diatomic sulfur. The equilibrium constant for the reaction is 2.27x10-3. 2H2S (g)  2H2 (g) + S2 (g) What is the concentration of hydrogen gas, [H2 ] , if [S2 ] =0.0540mol/L and [H2S] = 0.184 mol/L? • Write equilibrium expression first

  12. CW p 988 #6-11

  13. SOLUBILITY PRODUCT CONSTANT, Ksp • Some ionic compounds dissolve in water while others barely dissolve. • On dissolving, all ionic compounds dissociate into ions. Ex. NaCl(s)  Na+ (aq) + Cl– (aq) Mg(OH)2(s)  Mg2+ (aq) + 2OH – (aq) • Ksp (solubility product constant) expression is the product of the concentrations of the dissolved ions, each raised to the power equal to the coefficient of the ion in the chemical equation. For NaCl : ksp=[Na+ ][Cl– ] For Mg(OH)2 : ksp= [Mg2+ ][OH –]2

  14. Using solubility product constant • Help determine solubility of a compound (how much will dissolve) • Ex. AgI(s)  Ag+ (aq) + I – (aq) ksp=[Ag+ ][I – ] ksp= 8.5x10-17 at 298 K What is the solubility of AgI? (for every mole of AgI that dissolves, an equal number of Ag+ ions and I – ions forms.) [Ag+ ][I – ] = 8.5x10-17 (s)(s)= 8.5x10-17

  15. CW p 616 #20-21

  16. Calculating Ksp from solubility • What is the Ksp value for Ca3(PO4)2 at 298K if the concentrations at equilibrium are 3.42x10-7M for Ca2+ ions and 2.28x10-7M for PO43- ions?

  17. Calculating ion concentration • Ex. Magnesium hydroxide, Mg(OH)2 is used in the formulation of many medications, in particular those used to neutralize excess stomach acid. Determine the hydroxide ion concentration [OH-] if Ksp=5.6x10-12.

  18. Ex. The Ksp for silver carbonate, Ag2CO3, is 8.4x10-12 at 298K. What is the concentration of the silver ions, Ag+ ions?

  19. CW p617 #22,23

More Related