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Transcript of IV. Equilibrium - LaurenHill IV. Equilibrium 74 C. Reversible Reactions: Equilibrium Equilibrium...

  • IV. Equilibrium

    73

    13. The Nature of Chemical Equilibrium A. The Irreversible Reaction

    Most of the reactions we have considered so far have been irreversible, as implied by the one-way arrow: . In such a reaction, none of the product molecules reacted again to produce the original molecules.

    A non-technical example:

    Chemical examples: B. Steady State

    An open system can be in a steady state if the input rate of a substance equals the output rate. An open system is one that loses a substance from one “opening” and then gets that same substance back from another source.

    Non-technical examples:

    Chemical Steady State Examples from the Natural World:

  • IV. Equilibrium

    74

    C. Reversible Reactions: Equilibrium

    Equilibrium occurs in a closed system. Once a reaction attains equilibrium, the rate of the forward reaction equals the rate of the reverse reaction. At the macroscopic level, the reaction seems to have stopped, but at the molecular level it continues in both directions.

    Non-technical examples: Chemical Equilibrium Examples: Example 1 2 NO2(g) N2O4(g) + heat

    Macroscopic observations:

    At t = 0, we have no N2O4 (invisible) and 2 moles/L of NO2. (red- brown gas). With time,

    the brown colour fades. While these changes are occurring, we have not yet reached equilibrium.

    At equilibrium, the brown colour stops fading, and the concentrations of the two gases remain constant.

    Microscopically, what's happening?

  • IV. Equilibrium

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    Example 2 A large crystal of iodine is dropped into CCl4. a) Describe what you would see before equilibrium is reached and after equilibrium is

    attained. b) Also explain what occurs at the molecular level once equilibrium is established.

  • IV. Equilibrium

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    Exercises 1. Classify as one of the following: (1) Irreversible Reaction (2) Steady State (3) Reversible Reactions: Equilibrium a) water in an open dish left to evaporate b) 20 ml of alcohol in a sealed 500 ml jar c) water flowing into a pool at a rate of 2 L/s and leaving the pool at that rate. d) adding Mg to acid e) a saturated salt solution in a closed vessel 2. a) Sugar is added to a cup of coffee until no more sugar can dissolve. The top is sealed

    so that no coffee can evaporate, and the temperature is kept constant. Hey, we have equilibrium! Agree? b) Macroscopically, what would you see? c) Microscopically what is still going on?

    3. By burning a log, do you establish equilibrium eventually? Explain. 4. A single drop of water placed in a closed bottle may or may not establish a state of

    equilibrium between gas and liquid according to:

    H2O(l) H2O (g)

    Explain. What will affect whether equilibrium is established?

    5. Given: H2(g) + Cl2(g) 2 HCl(g) + 156 kJ

    Which information can you obtain from the above chemical equation? (yes/no) a) the ratio with which hydrogen and chlorine react ? b) the concentrations of the gases at equilibrium ? c) whether it is an equilibrium reaction? d) how fast equilibrium is reached? e) the physical state of the products and reactants? f) whether the reaction is exothermic? g) the reaction mechanism? 6. Give 1 everyday example for each of the following. Think of something that is not

    directly from your notes. (1) an irreversible reaction (2) Steady State (3) Reversible Reactions: Equilibrium

  • IV. Equilibrium

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    14. Disturbing Chemical Equilibrium: Le Chatelier's Principle

    In 1888, LeChatelier gave a succinct statement of the principle he had announced in 1884. It is:

    Every change of one of the factors of an equilibrium causes a rearrangement of the system in such a direction that the factor in question experiences a change in a sense opposite to the original change.

    Here is a different way of explaining the same concept.

    To predict how a system in equilibrium will react to a disturbance, view the forward and reverse reactions in competition with one another. The prevailing reaction will be the one that is getting what it needs from the disturbance. If one reaction is being more hampered than the other, obviously it will not win out.

    In Class Examples

    Each of the following reactions has reached equilibrium. What will be the effect on the equilibrium concentration of each substance when the change described is made?

    Example 1 2 H2(g) + 2 NO(g) N2(g) + 2 H2O(l)

    reacting hydrogen with a metal

    Example 2 SO2(g) + 0.5 O2(g) SO3(g) increasing the pressure on the system

    Henri LeChatelier (1850-1936)

  • IV. Equilibrium

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    Example 3 H2O(g) H2O(l) + heat

    (1) cooling the system

    (2) decreasing the pressure Example 4 Consider the following reaction:

    CaCO3(s) CaO(s) + CO2(g)

    H = (+) or endothermic limestone lime

    How would you maximize the amount of CaO (lime) produced?

  • IV. Equilibrium

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    Exercises 1. Each of the following reactions has come to equilibrium. But then you disturb the

    equilibrium. Predict the effect on the concentration of each other substance involved when the change described is made.

    a) 2 H2(g) + 2 NO(g) N2(g) + 2 H2O(g) removing water with a drying agent

    b) SO2(g) + 0.5 O2(g) SO3(g) using air instead of pure oxygen

    c) P4(s) + 6 H2(g) 4PH3(s) adding more H2

    d) FeO(s) + CO(g) Fe(s) + CO2(g) adding more carbon dioxide 2. In 1d, how could you decrease the amount of both FeO and CO while increasing the

    amount of CO2 present? 3. Will the reactants be favoured (will you get less product) if pressure is increased?

    Answer for each case.

    a) 2 H2(g) + 2 NO(g) N2(g) + 2 H2O (g)

    b) FeO(s) + CO(g) Fe(s) + CO2(g) 4. Water is boiling in a pressure cooker. The temperature is at 100 C. You would like to

    preserve some beans at a higher temperature, so you would like to prevent the water from boiling. You want to favour the reverse reaction, in other words.

    H2O(l) H2O(g) Do you raise or lower the pressure? Explain.

    5. When an NO2 and N2O4 mixture is placed in a syringe at room temperature,

    the following is observed: As the piston is gradually pushed in, the red brown colour of NO2 first darkens but then it gets progressively lighter. Explain.

    N2O4(g) 2 NO2(g) colourless red-brown

  • IV. Equilibrium

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    6. a) What will happen to the amount of SO3 if temperature is increased, given:

    SO2(g) + 0.5 O2(g) SO3(g) + 93 kJ ? b) In what direction will equilibrium “shift” if temperature is lowered for

    N2O4(g) 2 NO2(g) ? c) What colour change will you observe in the above as you lower temperature? (see #5 for colours) d) How would you change the temperature if you wanted to produce more NO2, given:

    2 NO(g) + O2(g) 2 NO2(g) -117 kJ

    7. A good humidity indicator can be made by coating paper with CoCl2 and observing its colour.

    Heat + [Co(H2O)6]Cl2(s) [Co(H2O)4]Cl2 (s) +2 H2O(g) Pink Blue a) If you were to observe a pink colour quickly forming at room temperature, what you

    conclude about the level of humidity? b) Suppose you were using the blue substance to gradually absorb moisture from your

    camera case. What would you do once the substance turned pink and stopped absorbing water? (How would you recycle the pink substance?)

    8. Anthocyanins are pigments that are responsible for most of nature’s blue, red and purple

    colours. Purple cabbage contains such pigments.

    Use these equations to answer the questions that follow.

    X + H+1 Y Blue red

    X + OH-1 Z Blue yellow

    a) While the purple cabbage is alive, which two pigments are present in its cells?

    b) When purple cabbage is cooked, the vacuoles (microscopic bags that contain pigment and acid ) burst, and the acid is neutralized by the rest of the cell’s alkaline contents.

    What colour will you see?

    c) If a blue anthocyanin solution is added to vinegar, what colour will it become?

    d) What two colours will you see if you add base(OH-1 ), one drop at a time, to a blue solution?

    e) How do you turn a yellow solution into a purple one?

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    15. Law of Chemical Equilibrium Since a reaction at equilibrium has fixed concentrations of products and reactants, we can calculate a constant at a given temperature.

    For aB + cD eF + gH

    Kc is known as the equilibrium constant. As long as temperature

    remains constant, the value of K will not change regardless of the initial

    amounts of reactants used.

    Keep in mind: You only include concentrations of aqueous and gaseous reactants/ products, not those of liquids or solids. Liquid and solid concentrations remain constant and so they already are imbedded in K. Exam