# Acid-Base Reactions, cont.

Consider the reaction of a weak acid and strong base: 50.0 mL of 0.25 M acetic acid is reacted with 18.0 mL of 0.50 M sodium hydroxide. Find the number of moles of each of the reactants and products after reaction.

• First recall that for the reaction of a weak acid and a strong base the reaction will go until one of the reactants is completely consumed.
• Writing the net ionic reaction we get (note that since it is a weak acid, we write it in the undissociated state):

CH3COOH + OH- H2O + CH3COO-

• First find moles of each reactant:
• acid = (50.0 mL)(1 L/1000 mL)(0.25 mole/L) = 1.25 x 10-2moles
• base = (18.0 mL)(1 L/1000 mL)(0.50 mole/L) = 0.900 x 10-2moles
• After reaction all of the base is consumed, so:
• base = 0
• acetate = 0.900 x 10-2moles
• acid = 1.25 x 10-2moles - 0.900 x 10-2moles = 0.35 x 10-2moles
• Water synthesized = 0.900 x 10-2moles.

## # Oxidation Numbers

For simple elemental ions it is easy to determine the charge on an atom, but in many other circumstances this is not the case. In order to name compounds and understand reactions we frequently need this information which is obtained from oxidation numbers.

Oxidation numbers are in essence an electronic accounting method in which electrons are assigned to a particular atom in a bond or interaction. As such they give an approximate picture of where electrons actually reside in compounds. We will find this information very useful later when we look at particular types of chemical reactions. Oxidation numbers are essential for nomenclature.

#### Oxidation numbers are most readily assigned using a simple set of rules:

1. In the formula for any substance the sum of the oxidation numbers of all the atoms in the formula is equal to the charge shown. Thus:
• For elements, such as Ar, O2, S8, etc. in the uncombined state the oxidation number for each atom must be 0, since no charge is shown and the atoms are equal to each other.
• For monoatomic ions the oxidation number equals the charge.
• For a compound the sum of the oxidation numbers of the atoms equals 0.
• For a polyatomic ion the sum of the oxidation numbers of the atoms equals the charge on the ion.
2. In compounds fluorine is always assigned an oxidation number of -1.
3. Alkali metals in compounds will always (for our class) be assigned an oxidation number of +1.
4. Alkaline-earth metals in compounds will always (for our class) be assigned an oxidation number of +2
5. In compounds oxygen is usually assigned an oxidation number of -2.
• Exception 1: in peroxides it is -1 while in superoxides it is -1/2. These will generally be obvious due to other rules (or the names).
• Exception 2: in combination with fluorine oxygen can be positive due to Rule 2 above, thus for OF2 oxygen is assigned an oxidation number of +2.
6. In compounds hydrogen is usually assigned an oxidation number of +1
• Exception: in metallic hydrides hydrogen is assigned an oxidation number of -1. These exceptions will be fairly obvious: NaH, CaH2, etc.
7. Aluminum will always (for our class) be assigned an oxidation number of +3, other elements in this Group will usually be assigned an oxidation number of +3.

#### Let's try the oxidation number rules on some examples:

• H2O2: Rule 1 & 6 exception (H cannot be more + than +1) H = +1, O = -1
• KO2: Rule 1, 3 & 5 exception 1 K = +1, O = -1/2
• OH-: Rule 1 & 5 -2 + H = -1, H = +1 - (-2) = -1
• MgH2: Rule 1 & 3 +2 + 2H = 0, 2H = -2, H = -1
• C2H3O2-: Rule 1 & 5 & 6 2C + 3(+1) + 2(-2) = -1, 2C = -1 - (+3) - (-4) = 0, C = 0
• MnO2: Rule 1 & 5 Mn + 2(-2) = 0, Mn = - (-4) = +4

#### Additional practice examples are available on the Study Module

Finally, note that in writing formulae, the element with the more positive oxidation number comes first. There are, of course, a few exceptions, the most well known being ammonia: NH3 (by the rules it should be H3N).

# Gases

Gases: Briefly discussed overall properties of gases (fills container, compressible, lo density, lo viscosity).

• Gas particles exert pressure.
• implication - gas particles have momentum (mass and velocity).

What is Pressure? Pressure is the force/unit area. Due to collisions of particle with walls of container etc.

Units of Pressure:

• mmHg - based on manometers. Two types:
• open tube - measures pressure relative to current atmospheric pressure.
• closed tube - measures pressure relative to contents of the enclosed volume at the closed end. (A barometer is an example where the enclosed space is "empty", that is it contains a vacuum since the vapor pressure of mercury is very low.)

• atm = 760 mmHg at 1 gravity = 1.01 Bar
• others include: psi (pounds/square inch), pascal, Torr (= 1 mmHg), millibar, etc.

# Gas Laws

Gas Laws describe the relationships between the four properties characterizing any gas:

• Amount of substance, (in moles)
• Volume, V (in Liters)
• Pressure, P (in atm, though often measured in mmHg)
• Temperature (in K) Note that Kelvins represent an "absolute" temperature scale with 0 at the lowest possible temperature and degrees the same size as °C. So 0 K = –273.15°C.

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