Unit 2: EquilibriumIcons are used to prioritize notes in this section. Make some notes: There are SOME important items on this page that should be copied into your notebook or highlighted in your printed notes. Copy as we go: There are sample problems on this page. I EXPECT YOU to copy the solutions into a notebook— Even if you have downloaded or printed the notes!!! Look at This: There are diagrams or charts on this page. Look at them and make sure you understand them. (but you don’t need to copy them) Extra Information: This page contains background information that you should read, but you don’t need to copy it. Review: There is review material on this page. It is up to you to decide if you want to make notes or highlight it, depending on how well you remember it. R
Unit 4(formerly Module 5) Equilibrium
Overview: Equilibrium is the concept of a system remaining “in balance”. A system in equilibrium does not change at the macroscopic level (the level that we can detect with our senses). True equilibrium should not be confused with homeostasis or “steady-state”, a process by which living organisms attempt to maintain a consistent internal conditions by absorbing or excreting materials from and to their environment. 11.0 Equilibrium
Equilibrium vs. Homeostasis • Equilibrium usually exists in a closed system, where materials cannot easily enter or leave. • Examples: • A reversible chemical reaction occurs inside a closed container. The reaction appears to have stopped, but at a molecular level changes are still going on. • A liquid in a sealed bottle does not appear to evaporate. • Homeostasis usually exists in an open system, where materials can enter and leave. • Examples: • A dog lives in a kennel. It eats and excretes roughly equal amounts, and therefore maintains a steady weight and internal conditions. • A cell in a human body maintains a balance of nutrients.
The Temporary Nature of Equilibrium • Although we study equilibrium as if it is an unchanging state, in reality dynamic processes as well as static forces may act upon the equilibrium. • Eventually something will upset the equilibrium, temporarily or permanently throwing it “out of balance” • Often a new equilibrium will be re-established after the original equilibrium is upset.
Examples • A precariously balanced rock formation can endure in “equilibrium” for centuries. Suddenly the balanced rock falls, temporarily disturbing the equilibrium. • A reversible chemical reaction inside a beaker has reached a state of equilibrium and appears to have stopped. A researcher adds more of one of the reactants, upsetting the equilibrium. The reaction temporarily resumes until a new equilibrium is established.
The Old Man of the Mountain For two hundred years a precarious rock formation in New Hampshire was said to resemble the face of an old man. It had become a symbol of New Hampshire, appearing on postcards, road signs and coins On May 3, 2003 The rock face collapsed. In terms of equilibrium, we could say that this was a static equilibrium that endured for centuries, until it was disturbed by a spring storm. Afterwards a new equilibrium was established, that unfortunately no longer resembled a face.
Chapter 11 Qualitative Aspects of Chemical Equilibrium
What is an equilibrium? What properties does it have? How can we distinguish dynamic, and static equilibria and tell them apart from simple steady states? 11.1 Qualitative Aspects of Equilibrium
Static Equilibrium • Static equilibrium exists when a system remains unchanging without any active, dynamic processes involved • One rock sitting on another is at static equilibrium, even if it is balanced precariously. No dynamic processes are acting on it, so it remains unchanged. • Static equilibrium is a bit boring. We seldom deal with it in chemistry. Note: Gravity is not considered to be a dynamic force. It is a static force.
Dynamic Equilibrium • Dynamic equilibrium is the result of two opposing, active processes occurring at the same rate. No visible changes take place, but there are constant changes in the particles at a microscopic level. • There are several types of dynamic equilibrium of interest to chemists, which are described on the following slides including: • Phase Equilibrium • Solubility Equilibrium • Chemical Equilibrium
A dynamic equilibrium is a bit like a hockey game. • Barring penalties, there is always the same number of players on the ice, but some players are constantly leaving the bench as others return to it. 19 11
Phase Equilibrium(1st type of dynamic equilibrium) • Phase equilibrium is a dynamic equilibrium that occurs when a single substance is found in several phases or states within a system as the result of a physical change. • Example: • In a closed bottle a water may exist as both a liquid and a gas at the same time (eg. Water vapour above liquid water). As water molecules evaporate from the liquid phase, other water molecules condense from the gaseous phase.
Solubility Equilibrium(2nd type of dynamic equilibrium) Solubility Equilibrium occurs when a solute is dissolved in a solvent, and an excess of the solute is in contact with the saturated solution. • Example: If you add too much sugar to a cup of tea, the tea becomes saturated with sugar and no more appears to dissolve. In fact, some molecules of sugar are dissolving as other molecules recrystallize back into solid sugar.
Chemical Equilibrium(3rd type of dynamic equilibrium) • Chemical equilibrium occurs when two opposing chemical reactions occur at the same rate, leaving the composition of the system unchanged. • Example: Dinitrogen tetroxide (N2O4) and nitrogen dioxide (NO2) can exist in the same container. Each can change into the other, and at equilibrium they do so at the same rate. N2O4 2 NO2 This is the most important type of equilibrium in chemistry!!
Some chemical reactions are easily reversed, like the electrolysis of water. Others, such as the burning of wood are impossible to reverse under laboratory conditions. 11.2 Irreversible and Reversible Reactions
Irreversible Chemical Reactions An irreversible reaction is a reaction that can only occur in one direction, from reactants to products. • This definition is assumed to refer to reactions occurring under normal laboratory conditions. • In a lab, it is easy to burn a piece of wood. It is impossible, under laboratory conditions, to turn ash, smoke, carbon dioxide and water back into wood. • The growth of a tree does allow wood to be produced from materials that might include wood ashes, but growing a tree takes decades, and requires countless changes involving many complex chemical mechanisms and multiple organic catalysts (enzyme systems). Burning is therefore NOT considered reversible!
Remember!Irreversible = One Way Reactants Products Reactants can become products, but products cannot turn back into reactants.
Reversible Chemical Reactions • Some reactions are easily reversed using common laboratory procedures. • For example, it is possible to decompose water into hydrogen and oxygen in a electrolytic cell, and equally possible to synthesize water from hydrogen and oxygen in a fuel cell: 2H2O + electrical energy 2 H2 + O2 2H2 +O2 2H2O + electrical energy
Remember!Reversible = Both Ways Reactants Products Reactants can become products, and products can also turn back into reactants. Also show as: reactants products
Reversibility and Equilibrium Only reversible reactions can produce a true dynamic chemical EQUILIBRIUM
A Chemical System is at Equilibrium if it meets these Criteria: • The System is Closed, for example by being sealed inside a container so material cannot enter or leave. • The Change is reversible. The reaction or change can proceed in both direct and reverse directions. • There is no Macroscopic Activity. Nothing seems to be happening – the properties of the system are constant. These unchanging properties can include: colour, amount of undissolved solute, concentration, pressure etc. • There is Molecular Activity. Reactions continue at the microscopic or molecular level
Definitionsof some easily confused terms • Macroscopic: Occurring at the level we can detect with our senses, as opposed to microscopic. Observable changes. • Microscopic: Occurring at a level below what we can see. Too small to observe without instruments. • Dynamic Equilibrium: A balance that involves two opposing active processes that are occurring at the same rate. This contrasts with Static Equilibrium and Steady State (Homeostasis). • Static Equilibrium: A balance that does not involve active processes. • Steady State (including Homeostasis): An apparent balance that occurs in an open system. An unchanging set of properties is maintained, but materials enter and leave the system.
Page 287 • Read all the questions, make sure you understand them, be prepared to answer them verbally next class.
Henri Louis Le Châtelier (1850-1937) was a French chemist who is most famous for his studies of chemical equilibrium. In addition he studied metal alloys and, with his father, was involved in the development of methods of purifying aluminum 11.4 Le Châtelier’s Principle
Equilibrium is not eternal • An equilibrium can exist for a long time, only to change (be upset) when certain conditions change. • After it is upset, there is a period of adjustment, then a new equilibrium is established. New Equilibrium Original Equilibrium Adjusting... Equilibrium Upset
Some factors that might upset an equilibrium. • Which of these factors do you think might affect the amount of reactant and product at equilibrium? • Maybe Temperature? • Maybe Pressure? • Maybe Concentration of reactants and products? • Maybe Catalyst? • Most of them do, but one does not. • We’ll see which one doesn’t a bit later!
Henri LeChâtelierstudied many of these factors to see how they could effect a system at equilibrium. • He found some factors could favour the direct reaction, increasing the amount of product. • Others could favour the reverse reaction, increasing the amount of reactant. • Regardless, eventually an equilibrium was re-established, but with new amounts of product and reactant.
When an Equilibrium is Upset... • Henri LeChatelier stated the following generalization: • “If the conditions of a system in equilibrium change, the system will react to partially oppose this change until it attains a new state of equilibrium.” • In other words: The system will establish a new equilibrium. .
Will any reversible reaction will eventually reach equilibrium? • Answer: yes, as long as it in a closed system. • Will the amount of product always equal the amount of reactant at equilibrium? • Short answer: NO! they are not always equal. • At equilibrium the amount of “reactant” and “product” may vary depending on several factors.
The Effect of a Catalyst • Adding a catalyst to a system already at equilibrium will have NO EFFECT. • Adding a catalyst to a system that has not yet reached equilibrium will cause it to reach equilibrium faster. • Why? A catalyst increases both forward and backward rates equally, so the final result will be the same, but the process of reaching equilibrium will be faster.
Effect of Temperature • Increasing the temperature will favour the endothermicreaction. • Decreasing the temperature will favour the exothermic reaction.
Effect of Pressure • Increasingthe pressure may favour the reaction that produces fewer gas particles • Decreasing the pressure may favour the reaction that produces more gas particles. • Note: only reactants or products that are in the gaseous state are counted towards the effects of pressure.
Effects of Concentration • The effects of concentration of the reactants and products are most important of all. Increasing or decreasing the concentration of Reactants WILL have an effect. • Remember: Pure solids(s)and liquids(l) do NOT have a variable concentration. • Before we can see the effects of changing a concentration we should remember what LeChatelier said...
Remember LeChatelier’s Principle: • “If the conditions of a system in equilibrium change, the system will react to partially oppose this change until it attains a new state of equilibrium.” • In other words, any “stress” or change that you make to the system will cause it to react in a way that tries to (partially) undo the change that you made.
Changes in Concentration • If you increase the concentration of a reactant (gas or aqueous),the equilibrium will shift to use up some of the reactant you added. • If you increase a product (gas or aqueous), the equilibrium will shift to reduce the product and make more reactant H2(g) + I2(g) 2HI(g) If you increase the concentration of Hydrogen… …The system will react to reduce the amount of Hydrogen… …by shifting the reaction towards the product
H2 + I2 2HI Sudden increase in[H2]. • The equilibrium changes: • [HI] goes way up, • [I2] goes down. • The equilibrium changes: • [HI] goes way up, • [I2] goes down. This causes [H2] to adjust towards its original level
Note: • Adding more of a undissolvable solid or pure liquid normally has NO EFFECT on a system at equilibrium. • It can only affect the equilibrium if the concentration changes, and usually only gases and aqueous solutions have variable concentration. • However: • adding water to an aqueous solution can change its concentration by dilution. • Adding solid to a solution that is not saturated MIGHT increase its concentration (if it dissolves!).
Changes in Temperature • Increasing the temperature causes the equilibrium to shift in the direction that absorbs heat (the endothermic direction). • Decreasing the temperature shifts the equilibrium in the exothermic direction. 2SO2 + O2 2SO3 + Heat If you increase the temperature… …The system will react to reduce the temperature… …by shifting the reaction towards endothermic side
2SO2 + O2 2SO3 + heat Sudden increase in temperature. The endothermic reaction kicks in, getting rid of some SO3, and creating more SO2 and more O2 and cooling things off This causes a lowering of the temperature, moving it towards its original level
Changes in Pressure(only affects systems where one or more materials are gases) • Increasing the pressure causes the equilibrium to shift in the direction that has the fewest gas molecules. • Decreasing the pressure shifts the equilibrium in the direction that produces more gas molecules. 8 molecules 4 molecules If you increase the pressure… …the system will create fewer molecules… …to reduce the pressure.
N2 + 3 H2 2 NH3 Sudden increase in pressure. Pressure partially adjusts towards the original level.. When the pressure increases, the equilibrium adjusts, making more NH3 (since 2 molecules NH3 are fewer particles than 3 molecules of H2 plus 1 molecule of N2)
Special Note Re. Gases • Only gases can be affected by pressure. • If only one side of an equation has gases, then... • Increasing pressure will favour the side with no gases. • Decreasing the pressure will favour the side with gases. I2(s) I2(g) Decreased pressure Increased pressure
Chapter 12 The Quantitative Aspects of Equilibrium
In this section we will explore the mathematical aspects of equilibrium, including: The Equilibrium Constant (Kc) The Equilibrium Law 12.1 Chapter 12
The Equilibrium Constantand Equilibrium Law Expressions • The equilibrium constant (Kc) is: • A number • derived from an equilibrium law expression. • a ratio between the concentration of products and the concentration of reactants of a reversible reaction at equilibrium, but… • With each aqueous or gaseous product and reactant raised to the power of its corresponding coefficient
Deriving the Equilibrium Law Warning: Math Content Ahead • The next two slides show how the equilibrium law was derived from the rate law. If you want to understand the relationship between rates and equilibrium you should follow this. • If you just want to use the equilibrium law to find Kc, you can skip forward three slides.
aA + bB↔cC + dD reactantsproducts Forward rate: rdir = kdir [A]a [B]b Reverse rate: rrev = krev [C]c [D]d At equilibrium, rdir = rrev so, through the magic of algebra… A Generalized Reversible Reaction with reactants and products Forward reaction = dir Reverse reaction = rev Products on top Reactants on bottom