Chapter 7:Chemical Reactions: Energy, Rates, and Equilibrium

• There are two fundamental kinds of energy: potential and kinetic.

  • Potential energy is ==stored energy==. The water in a reservoir behind a dam, an automobile poised to coast downhill, and a coiled spring have potential energy waiting to be released.
  • Kinetic energy, is the ==energy of motion==. When the water falls over the dam and turns a turbine, when the car rolls downhill, or when the spring uncoils and makes the hands on a clock move, the potential energy in each is converted to kinetic energy.

• In chemical compounds, ==the attractive forces between ions or atoms are a form of potential energy==.

• When these attractive forces result in the formation of ionic or covalent bonds between ions or atoms, ==the potential energy is often converted into heat—a measure of the kinetic energy== of the particles that make up the molecule. Breaking these bonds requires an input of energy.

• The term “stable” is used in chemistry to describe a substance that has little remaining potential energy and consequently little tendency to undergo further change.

• Bond dissociation energy is defined as the ==amount of energy that must be absorbed to break the bond and separate the atoms== in an isolated gaseous molecule.

• The greater the bond dissociation energy, the more stable the chemical bond between the atoms or ions.

• A chemical change that absorbs heat, like the breaking of bonds, is endothermic.

• The reverse of bond breaking is bond formation, a process that releases heat and is exothermic.

• For endothermic processes, heat is absorbed (gained) and is indicated by a positive sign. For exothermic processes, heat is released (lost) and is indicated with a negative sign.

• The difference between the heat energy absorbed in breaking bonds and the heat energy released in forming bonds is called the heat of reaction and is a quantity that we can measure.

• When the total strength of the bonds formed in the products is greater than the total strength of the bonds broken in the reactants, the net result is that energy is released and the reaction is exothermic.

• When the total energy released upon bond formation in the products is less than the total energy added to break the bonds in the reactants, the net result is that energy is absorbed and the reaction is endothermic.

• Important conclusions about heat transfer:

  • An exothermic reaction releases heat to the surroundings; ∆H is negative.
  • An endothermic reaction absorbs heat from the surroundings; ∆H is positive.
  • The reverse of an exothermic reaction is endothermic.
  • The reverse of an endothermic reaction is exothermic.
  • The amount of heat absorbed or released in the reverse of a reaction is equal to that released or absorbed in the forward reaction, but ∆H has the opposite sign

• A spontaneous process is one that, once started, proceeds on its own without any external influence.

• A nonspontaneous process, by contrast, takes place only in the presence of a continuous external influence. Energy must be continually expended.

• The amount of disorder in a system is called the system’s entropy, symbolized by S and expressed in units of Joules per mole-kelvin.

  • The greater the disorder, or randomness, of the particles in a substance or mixture, the larger the value of S.
  • ==Gases have more disorder and therefore higher entropy== than liquids because particles in the gas move around more freely than particles in the liquid. Similarly, liquids have higher entropy than solids.

• The entropy change (∆S) for a process has a positive value if disorder increases because the process adds disorder to the system. The melting of ice to give water is an example. Conversely, ∆S has a negative value if the disorder of a system decreases.

• Two factors determine the spontaneity of a chemical or physical change: the release or absorption of heat, ∆H, and the increase or decrease in entropy, ∆S.

• The value of the free-energy change, ∆G, determines spontaneity. A negative value for ∆G means that free energy is released and the reaction or process is spontaneous. Such events are said to be exergonic. A positive value for ∆G means that free energy must be added and the process is nonspontaneous. Such events are said to be endergonic.

  ∆G =∆H-T∆S.

  • A spontaneous process, once begun, proceeds without any external assistance and is exergonic; that is, free energy is released and it has a negative value of ∆G.
  • A nonspontaneous process requires continuous external influence and is endergonic; that is, free energy is added and it has a positive value of ∆G.
  • The value of ∆G for the reverse of a reaction is numerically equal to the value of ∆G for the forward reaction but has the opposite sign.
  • Some nonspontaneous processes become spontaneous with a change in temperature

• The amount of energy the colliding particles must have for productive collisions to occur, an amount called the activation energy (Eact) of the reaction. The size of the activation energy determines the reaction rate, or how fast the reaction occurs.

• Several things can be done to help reactants over an activation energy barrier and thereby speed up a reaction:

  • Temperature:

  one way to increase reaction rate is to add energy to the reactants by raising the temperature. With more energy in the system, the reactants move faster, so the frequency of collisions increases.

  • Concentration:

  another way to speed up reaction is to increase the concentrations of the reactants. As the concentration increases, reactants are crowded together, and collisions between reactant molecules become more frequent.

  • Catalysts:

  a third way to speed up reactions is to add a catalyst - a substance that accelerates the rate of a reaction without changing itself during the process.

• Reactions which easily go in either direction, are called reversible reactions and are indicated by a double arrow in equations. The reaction read from left to right as written is referred to as the forward reaction, and the reaction from right to left is the reverse reaction.

• No matter which pair of reactants is mixed together, both reactions occur until ultimately the concentrations of reactants and products reach constant values and undergo no further change. At this point, the reaction is in a state of chemical equilibrium.

• Chemical equilibrium is an active, dynamic condition. All substances present are continuously being made and unmade at the same rate, so their concentrations are constant at equilibrium.

• For a reversible reaction, then, the rates of both the forward and the reverse reactions must depend on the concentration of reactants and products, respectively.

  • When a reaction reaches equilibrium, the rates of the forward and reverse reactions are equal, and the concentrations of reactants and products remain constant.

• The equilibrium constant K is the number obtained by multiplying the equilibrium concentrations of the products and dividing by the equilibrium concentrations of the reactants, with the concentration of each substance raised to a power equal to its coefficient in the balanced equation.

  • K much smaller than 0.001: Only reactants are present at equilibrium; essentially no reaction occurs.
  • K between 0.001 and 1 :More reactants than products are present at equilibrium.
  • K between 1 and 1000 :More products than reactants are present at equilibrium.
  • K much larger than 1000 :Only products are present at equilibrium; reaction goes essentially to completion.

• Le Châtelier’s principle states that when a stress is applied to a system at equilibrium, the equilibrium shifts to relieve the stress.

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