Some compounds dissolve in water by dissociating into ions;
Examples include sodium chloride (NaCl), hydrochloric acid (HCl),
acetic acid or vinegar (CH3COOH), sodium hydroxide
(NaOH), and ethanol (CH3CH2OH)
To start the process of breaking apart the ions some energy is
required
This energy is more than regained in the process of
hydration
Hydration is a process in which opposite charges attract one
another; in this case the partial positive charge on the hydrogen
in H2O is attracted to anions and the partial negative
charge on the oxygen is attracted to cations
In hydration each ion is surrounded by a group of opposite charges,
thereby lowering the overall energy of the solution
Remember that these substances in their pure state are not
separated ions: they are electrically neutral and do not conduct
electricity
Conductivity
Molecular compounds such as pure acetic acid (vinegar without any
water at all) do not conduct electricity
Under the influence of water, in the process of hydration, these
compounds can be broken into ionic pieces; when this happens the
solution conducts electricity
Ionic compounds are also poor conductors of electricity when not in
solution
Hydration frees ions to move around in the solution; it is this
motion which enables an electric current to flow
Cations move toward the negatively charges anode; anions move
toward the positively charged cathode
Substances which conduct electricity when placed in solution are
called electrolytes
Strong electrolytes are substances that, when dissolved in
water, conduct electricity very well (example: NaCl); these
substances, for all intents and purposes, break up completely
into ions in solution
HCl(aq) + H2O(l) →
H3O+(aq) +
Cl–(aq) [H+ in solution is usually shown as H3O+]
Weak electrolytes are substances that, when dissolved in water,
conduct electricity only poorly (example: acetic acid); these
substances break up into ions to only a small extent in
solution
CH3COOH(aq) + H2O(l)
↔ CH3COO–(aq) +
H3O+(aq)
Nonelectrolytes are substances that do not conduct electricity
when dissolved in water (example: sugar); these substances do
not break up into ions in solution
C6H12O6(s) +
H2O(l) ↔
C6H12O6(aq) +
H2O(l)
Whether a substance is a strong electrolyte or a weak electrolyte
has less to do with the concentration of the solution and more to
do with the equilibrium of the hydration process
Reversible Reactions and Equilibrium
Hydration is an example of a reversible reaction; most chemical
reactions are to some extent reversible
Weak electrolytes, upon introduction into water, begin to
dissociate into ions but most of the particles remain in molecular
form
But the particles in molecular form are just as likely to
dissociate as molecules which have already become ions
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What is to prevent these other particles from dissociating?
Nothing; they do dissociate but at the same rate that ions are
combining to reform molecular particles
In this way the total number of ions and molecules stays the same
while each is constantly becoming the other
Strong electrolytes simply dissociate to a greater extent than weak
electrolytes; the equilibrium is here weighted toward the side of
the equation with the ions rather than the molecular particles so
that almost no molecular particles can be found in the solution
This state of affairs, in which the number of ion pairs and
molecular particles stays the same, is called equilibrium
Equilibrium should not be thought of as the end of a
reaction but rather as the point at which no net change
occurs
Think of this like filling a funnel with water; there is a certain
rate that you can pour the water in to exactly match the amount of
water flowing out: the level of water in the funnel doesn’t
change even as the water continues to flow
The best description of a reaction at equilibrium involves a comparison of the rates of the forward and reverse reactions; when these rates are equal, the reaction is at equilibrium
Equilibria of many different kinds exist:
Physical equilibrium can be reached between a solid salt and a saturated solution of that salt (NaCl(s) ↔ NaCl(aq))
Physical equilibrium can be reached between liquid water and gaseous water in a closed container (H2O(l) ↔ H2O(g))
Physical equilibrium can be reached between molecular particles and ions in solution (CH3COOH(aq) + H2O(l)
↔ CH3COO–(aq) +
H3O+(aq))
Chemical equilibrium can be reached between reactants and products (N2(g) + O2(g) ↔ 2NO(g))
Non-equilibrium Systems
Some reactions do have an end-point; these reactions are said to go
to completion
Examples of reactions that go to completion are those that result in
the formation of a gas (such as our experiment last week
with the production of hydrogen) or in the formation of a precipitate
These reactions are not considered equilibria because a product
forms that becomes unavailable to react in the reverse direction; a
precipitate is insoluble and will not react with ions in solution;
a gas escapes to the atmosphere and is no longer even physically
present to react
Le Châtelier’s Principle
Left undisturbed reversible reactions will reach equilibrium; a point at which no net reaction is occurring; or the reaction is progressing backward at the same rate it is progressing forward
Equilibrium can be disturbed however, causing the reaction to progress more in one direction than another until a new equilibrium is reached
There are many ways equilibrium can be disturbed, among them are:
A change in temperature; i.e., any addition or subtraction of energy from the system
A change in pressure
The removal or addition of a product or reactant
A clear statement from Mr. Le Châtelier himself: “Any change in one of the variables that determines the state of a system in equilibrium causes a shift in the position of equilibrium in a direction that tends to counteract the change in the variable under consideration.”
Taking the example of adding heat to a system, let’s look at two reactions: an exothermic reaction and an endothermic reaction
In an exothermic reaction heat is produced by the reaction: adding heat will cause the formation of reactants
In an endothermic reaction heat is consumed by the reaction: adding heat will cause the formation of more products