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Intro to reactions in solution

Reactions in solution 1


We are going to be looking at reactions in aqueous solutions (in water). The reason for this is that in order for a reaction to occur, the atoms, molecules and ions have to be able to interact with each other. They can do this if they are in either liquid or gaseous phase, but it is harder for this to happen if the compounds in question are in solid phase. While you can have solutions of solids, liquids or gases, but we will speak almost entirely on liquids. By convention, we call the thing that is present in greater concentration the "solvent" (this is the thing doing the dissolving) and the thing that is dissolved in it is called the "solute" (this is the thing dissolved).
BTW,
here are the pages from the white board…actually written last year.

Electrolytes:


The definition is functional: a solution that conducts electricity. In practice, this means that there are ions in solution, as opposed to molecules, since it is the presence of positive and negative ions that allows conductivity. Any electrolyte must involve some ions separating from each other. If a compound produces a lot of ions in solutions, such as when an ionic compound like NaCl dissolves, it is a strong electrolyte. Things that ionize only a little, like weak acids, will be weak conductors. Weak electrolytes, weak acids and weak conductors should be linked in your mind. A strong acid results in a low pH and will be good a conductor. Strong and weak bases will be analogous.

Solutions of ionic compounds in water are not "molecules" in solution, but were rather the separate ions of the compound. If an ionic compound dissolves, it does so by separating into its component ions which are "hydrated," or surrounded by water molecules. The partially negative oxygen crowds around the positive ion and the partially positive Hydrogens crowds around the negative ion. Solubility implies that the condition of hydration is lower-energy for that compound than the crystal. If the bonds between the ions is too strong, the compound will not be soluble, because the water cannot break the bonds.

What types of reactions can we follow?


There are a few we will consider:
  1. Precipitation reactions: In these, two solutions of ions are mixed. When the cation of one compound interacts with the anion of the other, the bonds of this new compound are too strong for the water to keep in solution, and the new solid precipitates.
  2. Acid-base reactions: in these, you mix the acid and base, which will result in a solution of intermediate pH. The product of these is generally water and some soluble salt.
  3. Oxidation/reduction reactions. These involve the trading of electrons, and are most notable in that they are the underlying chemistry of batteries and you.

Net ionic reactions:


We write equations usually as what we call balanced molecular reactions. For example

MgCl
2 (aq) + 2AgNO3 (aq) 2AgCl (s) + Mg(NO3)2 (aq)

In this notation, the small "s" in parentheses is "solid" and the "aq" is the aqueous form. But, we know that there are not molecules in solution, they are ions. So we can write the "complete ionic reaction:

Mg
2+ + 2Cl- + 2Ag+ + 2NO3+ 2AgCl(s) + Mg2+ + 2NO3+

Notice that we do not separate the AgCl because it is not aqueous. It is solid. When we look at this, the NO
3+ and the Mg2+ are the same on each side. By the rules of algebra, we can remove them. These ions are considered the "spectator ions," because they do not participate in the reaction.
After we remove the spectator ions, we have the net ionic reaction. That is the reaction written from the perspective of the things that actually change, is:
Ag
+ + Cl- AgCl (s)
You may note that I have removed the coefficient "2" from the net ionic reaction. That is the correct way to write it. The net ionic equation is often the one we will use because it is the simple expression of what happens.
Precipitation How do you know what precipitates and what does not? Here is a link to the solubility chart we use.
Okay, so some memorization may be necessary here. Wikipedia also has a
good chart. But, let's try to keep memorization to a minimum. Also, you will do some experiments that address this. You will look at solubility of various compounds, and find, for example, that all the salts of alkali metals are easily soluble. Also, that salts of the metals in the alkaline earth column are less soluble, but BaCl2 for example, is more soluble than CaCl2.

Acids/Bases


The first thing you should know is the definitions of acid and base. It turns out there are several. For our purposes, an acid is a proton (H
+) donor and bases are proton acceptors. For a "reaction" to happen, you need both an acid and a base. However, water is often the other player in the reaction because it can be both an acid and a base (it's just equally bad at both). That property makes it ideal as the solvent for the reactions we will consider.

pH.


You've heard the term; but, what does it mean? Consider the following reaction:


The double arrow means it can go either direction and, in this case, favors the leftward direction.
One of the water molecules on the left is the proton donor (the acid) and the other is the proton acceptor (the base). Water is, as I said, equally bad at both. So, water can spontaneously give up a proton while another accepts it to become OH
- and H3O+, but does so to a very low extent. How low an extent? Well, we could express that in terms of what the concentration of OH- and H3O+ are in distilled water. The concentrations should be equal, since you produce them in a 1:1 ratio. By the way, we will abbreviate "Concentration of" by using square brackets. [H3O+] means "concentration of H3O+.
The concentration of H
+ (or H3O+) in distilled water is 10-7. There is a mathematical operation, taking the "p" of something is to take the negative log. So, the pH is 7 in distilled water. Note, I could talk about the pOH just as easily. It also would be 7 in distilled water, since the [OH- ] is also 10-7.
The sum of pH + pOH always will equal 14. As pH goes down (increasing [H
+]) the pOH goes up by the same amount (decreasing the [OH-])

For a strong acid (one that dissociates fully in water), the concentration of H+ can be taken directly from the concentration of the acid. So, if know the [HCl] is 0.025M. You can calculate the pH as follows:
All of the HCl will dissociate into H
+ and Cl-. So, the [H3O+] will be 0.025M. Take the log of that and you get -1.6. Convert the sign and you get pH=1.6.

Conjugate acid/bases and Ka


What about weak acids? Well, weak acids, such as acetic acid (vinegar) don't dissociate fully.
The formula for acetic acid can be written:
CH
3COOH
That last "H" is the one that comes off as a proton. That's one reason we write it that way. In water
CH
3COOH +H2O CH3COO- + H3O+ The double arrow means the reaction can go both ways. So, reading the reaction left to right, CH3COOH is the proton donor (the acid) and H2O is the base (proton acceptor). Reading it right to left, the H3O+ is the acid and CH3COO- is the base. So, CH3COOH and CH3COO- are an acid and its conjugate base. Water and H3O+ are the base and conjugate acid of each other.

If I tell you that you have 0.025M CH
3COOH, you cannot conclude the pH is 1.6, as we did for the HCl, because "only a fraction" of the acid dissociates. The question is: "what is that fraction?" If I knew that, I could figure out the [H3O+] and therefore the pH.
Well, that fraction would be just the concentration of the dissociated ions divided by the amount that stays as the intact molecule. We call that fraction the K
a. K stands for "constant" (in German) and 'a' is for acid. It turns out that at any given temperature, an Acid has a measurable fraction that dissociates.
We write it like this:

At 25℃, the Ka for Acetic acid is 1.8x10-5
Use that to determine what the pH of solution of 0.025M of acetic acid is. Hint, since the concentration of H+ and that of CH
3COO- are the same, you can treat the top of the fraction as "x2"


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