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More About Spontaneous Reactions

A Redox Reaction Example

Demonstration: A piece of copper wire in a solution of 2% AgNO3.

Initially the copper is a shiny copper color and the solution is clear. In less than one hour the solution is light blue and the wire is covered with shiny silver needles. What happened?

Copper metal became copper ions in solution and silver ions became silver metal.

Cu(s) + Ag+(aq) ---> Cu2+(aq) + Ag(s)       (unbalanced)

The Cu(s) loses electrons to become Cu2+(aq) ions and the Ag+(aq) ions gain electrons to become Ag(s).

Reactions that involve the exchange of electrons are called reduction and oxidation (redox) reactions. When a chemical species loses electrons we say that it is oxidized, and when a chemical species gains electrons we say that it is reduced.

The Cu(s) loses electrons to be oxidized to Cu2+(aq).
The Ag+(aq) gain electrons to be reduced to Ag(s).

We observed that copper metal and silver ions undergo a spontaneous reaction to produce copper ions and silver metal.

Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s)

Will this reaction continue indefinitely?

At some point we reach equilibrium concentrations and the reaction stops. (Remember that microscopically reactants and products continue to interchange, but we say the reaction has stopped when the concentrations reach a steady state). When we reach equilibrium we write the reaction with arrows in both directions.

Cu(s) + 2 Ag+(aq) <--> Cu2+(aq) + 2 Ag(s)

Which of the following reactions are spontaneous?

H+(aq) + OH-(aq) ---> H2O

H2O ---> H+(aq) + OH-(aq)

The first reaction is spontaneous and the second is not. Water does not spontaneously decompose. It does reach the following equilibrium:

H2O <--> H+(aq) + OH-(aq)

We use the right arrow when we are talking about a reaction at non-equilibrium concentrations and the double arrows when a reaction has reached equilibrium. You will see the same reaction written both ways, the different notation indicates that we are considering either non-equilibrium or equilibrium concentrations of the reactants and products.

Reactants that are mixed in non-equilibrium concentrations have potential energy that drives the reaction towards equilibrium. This potential energy is the difference in free energy between the products and reactants, DeltaG. The reactants will react to achieve equilibrium and reach a lower energy condition.

DeltaG = DeltaGo + RTln(Q)

where
DeltaGo is the free energy of reaction at standard concentrations
R is the gas constant, 8.3145 J/mol·K
T is absolute temperature
Q is the reaction quotient

This relationship shows that reactants at concentrations far from equilibrium will have greater free energy than reactants near equilibrium concentrations.

At equilibrium DeltaG = 0. We cannot extract energy from a system at equilibrium.

Also note that DeltaGo = -RTln(K)

That these two are related should make sense. K tells us relative concentrations of products to reactants at equilibrium, DeltaGo tells us how much energy we could get out of a reaction when they are at standard concentrations.


Don't Forget the Really Big Concepts in Chemistry

Kinetics describes how quickly or slowly a reaction occurs.

Thermodynamics describes the changes in the form of energy when a reaction occurs, for example, converting chemical energy to heat.

Equilibrium describes reactions in which the reactants and products coexist.


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