Chapter+12,+Section+1


 * This page contains information that is specific to Chapter twelve material. It is a good supplement to my lectures and to the textbook; different explanations and pictorial representations are offered. Take the time to look through the ideas and concepts presented here as we are discussing this chapter in class. Click here to see the procedure for the lab we will be performing for this Chapter 12 (Chemical Equilibrium).**

S11.A.1.1.2 Analyze and explain the accuracy of scientific facts, principles, theories,and laws.

If an aqueous solution conducts electricity, then it must contain ions. So measuring the conductance of solutions can tell you whether the solutes in the solution are dissociated into ions, and whether chemical reactions in solution are producing or consuming ions. Any solution, even one containing ions, provides considerable resistance to the flow of current through it. Conductivity is, roughly speaking, the reciprocal of this resistance -- high resistance means low conductivity; low resistance means high conductivity.



Circuit diagram of conductivity measurement. The solution completes a circuit that includes a battery of known voltage (left) and an ammeter (right). An internal computer chip converts the meter output to conductivity. A portable conductivity meter incorporates all these components into a single device.

Some Basic Principles of Conductivity
Conductivity is roughly proportional to the concentration of ions in solution, but all ions do not conduct equally. Ions that move through solution easily conduct better. For example, small, fast moving ions like hydrogen ion (H+ ) impart greater conductivity to solutions than do bulky ions like bromide ion (Br - ), or heavily hydrated ions like sulfate ion (SO42- ). Electrolytes are compounds that dissolve in water and dissociate, at least partially, into ions. In solution of elctrolytes, several different species might be present, including intact molecules and dissociated ions. Strong electrolytes dissociate completely into ions. For example, a 1.0 M solution of the ionic strong electrolyte AZ contains 1.0 M A+ ions, 1.0 M Z- ions, and 0.0 M AZ molecules. In other words, the ions are the only species present in a solution of a strong electrolyte. Diluting a 1.0-M solution of AZ with water to 0.50 M reduces the concentration of ions by one half, and thus reduces the conductivity of the solution by one half. Weak electrolytes dissociate incompletely. For example, a 1.0-M solution of the weak electrolyte BY might contain less than 0.10 M B+ ions, the same molarity of Y- ions (why the same?), and greater than 0.9 M BY molecules. In other words, in a solution of the weak electrolyte BY, the predominant species are BY molecules, while B+ and Y– ions are minor species (present, but not as numerous as the predominant species). The dissociation of BY is an equilibrium process, in which BY molecules are constantly dissociating (forward reaction) and reforming (reverse reaction), at identical rates: BY(aq) <==> B+(aq) + Y-(aq) In this reversible reaction, a B+ ion and a Y- ion form whenever a BY molecule happens to dissociate spontaneously (the forward reaction), while a BY molecule forms whenever a B+ ion and a Y- ion happen to collide (the reverse reaction). At equilibrium, the forward and reverse processes occur at the same rate, so BY dissociates and reforms at the same rate, and [BY], [B+], and [Y-] all remain constant. The probability of spontaneous dissociation of a BY molecule is constant, but the probability of a B+/Y- collision depends on the concentration -- low concentration makes collisions more rare, and makes the reverse reaction slower. This means that diluting a solution of BY slows down the reverse reaction more than it slows down the forward reaction. After dilution, dissociation outpaces collision until the concentrations of the ions rise, their collision rate increases, and the rate of the reverse reaction rises to match that of the forward reaction. As a result, dilution leads to dissociation of additional BY molecules before equilibrium is reached again. So diluting a 1.0-M solution BY to 0.50 M with water reduces the conductivity, but by less than the 50% expected with strong electrolytes.

1.3 Ionization & Dissociation
You should recall from our unit on solutions that electrolytic solutions are those that conduct electricity because the substance dissolves in water to produce ions. > This concept of strong and weak will have an important role in our discussions of acids and bases, so it is important to understand the difference now. Ionic compounds and some molecular compounds can produce electrolytic solutions, but a different name is //usually// given to the processes. ..
 * Solutions that conduct electricity well are **strong electrolytes** – they are good conductors because they break down well and produce many ions in solution.
 * **Weak electrolytes** do not conduct electricity as well because fewer ions are produced in solution.

**Dissociation**
When ionic compounds dissolve to produce ions the process is typically called **dissociation**. Dissociation of ionic compounds occurs when water molecules “pull apart” the ionic crystal. This occurs due to strong attractions between the polar ends of the water molecule and the positive and negative ions within the crystal. Water molecules then surround the positive cations and negative anions; this is called hydration. Examples of dissociation equations: What do you notice about the last three reactions shown above? All three of them produce hydroxide ions, OH–. These three compounds - NaOH, KOH, and Mg(OH)2 are all Arrhrenius bases since they produce hydroxide ions in solution.
 * NaCl(s) → Na+(aq) + Cl-(aq) ||
 * Na2SO4 (s) → 2 Na+(aq) + SO42-(aq) ||
 * (NH4)3PO4 (s) → 3 NH4+(aq) + PO43-(aq) ||
 * NaOH(s) → Na+(aq) + OH– (aq) ||
 * KOH(s) → K+(aq) + OH– (aq) ||
 * Mg(OH)2 (s) → Mg2+(aq) + 2 OH– (aq) ||

There are two important things to notice about writing dissociation equations: > > > An ion, such as the sodium ion Na+ is not the same as a sodium atom, Na. Be sure to get in the habit of writing charges for all ions; refer to your [|Table of Common Ions] when needed, but by now you should be memorizing the charges of ions we commonly use, including polyatomic ions.
 * 1) Generally DO NOT include H2O as a reactant. We know something has been dissolved in water when we see the (aq) notation. We will make some exceptions later to this rule, however.
 * 1) Ion charges **MUST BE** included!

**Ionization**
When molecular compounds dissolve in water to produce ions the process is typically called **ionization**.

Recall that molecular compounds are held together by covalent bonding; ionic compounds are held together by ionic bonding

Most molecular compounds do **not** undergo ionization. Acids are an exception. All acids produce hydrogen ions in solution. Some examples:


 * HCl (g)→ H+(aq) + Cl-(aq) ||
 * H2SO4 (g)→ 2 H+(aq) + SO42-(aq) ||
 * HC2H3O2 (l) || [[image:http://www.saskschools.ca/curr_content/chem30_05/graphics/site_wide/darrow.gif width="20" height="10" caption="in equilibrium with"]] || H+(aq) + C2H3O2-(aq) ||
 * (The reason a double arrow is used for acetic acid will be discussed later) ||

The reactions as written above are actually simplified versions of what really occurs. Evidence suggests that the hydrogen ion, H+, actually bonds to a water molecule (H2O) to form the **hydronium ion**, H3O+:

For our class, it won't make a difference if you write an acid ionization reactions as producing hydrogen ions (H+) or hydronium ions (H3O+). For our purposes they will represent the same thing. You should be comfortable using either; one will mean the same as the other.
 * HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq) ||
 * H2SO4 (g) + H2O(l)→ 2 H3O+(aq) + SO42-(aq) ||
 * HC2H3O2 (l) + H2O(l) || [[image:http://www.saskschools.ca/curr_content/chem30_05/graphics/site_wide/darrow.gif width="20" height="10" caption="in equilibrium with"]] || <span style="font-family: Arial,Helvetica,sans-serif;">H3O+(aq) + C2H3O2-(aq) ||