CHEM 1151K
October 19, 2004
Examples:
Page 125 problem 15c
Find the mass in grams of 0.250 moles of C6H12O6
First, calculate the molar mass in g/mol. Add the atomic weights of all atoms
shown in the chemical formula:
(6 x 12.022) + (12 x 1.008) + (6 x 15.9994) = 180.1584 g/mol
carbons
hydrogens
oxygens
1 mole C6H12O6 = 180.158 grams
Now, use this to make a conversion factor for moles  grams.
 180.158 grams 
0.250molC 6 H 12 O6 
  45.039 gramsC 6 H 12 O6
1mole


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Problem 24, page 125 correction: The formula for TNT should read: C7H5N3O6
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Problem 21 page 125
What mass of ammonia, in grams, can be made from 440 g of H2?
1 N2 + 3 H2  2 NH3
(balanced)
Use the flow-chart: grams A  moles A  moles B  grams B
Substance "A" is hydrogen, substance "B" is the ammonia.
Step 1: Convert grams of hydrogen to moles of hydrogen. You need the molar mass of
H2 to do this. If you double the atomic weight of hydrogen, you get:
1 mole H2 = 2.016 grams H2
Use this relationship as a conversion factor to convert the given amount of
hydrogen into moles of hydrogen.
 1molH 2
440 gH 2 
 2.016 gH 2

  218molH 2

Step 2: Convert the moles of hydrogen found in step 1 to moles of ammonia. To convert,
use the balancing coefficients from the balanced chemical reaction shown above. The
reaction tells us that we get 2 moles of ammonia for every 3 moles of hydrogen that are
reacted.
 2molNH 3 
  146molNH 3
218molH 2 
 3molH 2 
Step 3: Lastly, we convert the moles of ammonia found in step 2 into grams of ammonia.
We need to find the molar mass of ammonia to do this. Add up the atomic weights of 1
nitrogen and 3 hydrogens.
1 mole NH3 = 17.03 grams NH3
Again, apply this as a conversion factor, this time from moles  grams.
 17.03gNH 3
146molNH 3 
 1molNH 3

  2490 gNH 3

Yes, that's a lot of ammonia! Remember, you started out with 440 grams of hydrogen.
Plus remember that 1 mole of hydrogen gas weighs about 2 grams, but a mole of
ammonia is much heavier at 17 grams.
Now back to our discussion of the energies of reactions.
Review: Endothermic processes absorb heat from the surroundings. Exothermic
processes liberate heat to the surroundings.
See the energy diagrams at the end of the Oct. 14 lecture.
Whether a reaction is endothermic or exothermic it may need additional energy to get
started. This is called the activation energy. The combustion of gasoline is obviously
and exothermic process (gives off heat), but a bucket of gasoline is unlikely to just start
burning even though the air has plenty of oxygen to support combustion. However, one
little lit match will get the reaction going!!!! The match provides the initial activation
energy. On energy diagrams the activation energy appears as a "hill" between the
reactants and products. It's value is the difference between the peak energy of the hill
(transition state or TS) and the energy of the reactants.
This example is exothermic. Notice the energy of the reactants is greater than the
products. The excess energy can be released as heat.
Here's an interesting website for more help on this:
http://scidiv.bcc.ctc.edu/wv/7/0007-004-reac_rate2.htm
Rates of Reaction - the rate tells the speed of the reaction (fast or slow?)
It can be expressed as the amount of a product produced per unit time (like per minute),
or the amount of a reactant used up per unit time.
Some example units: g/sec mol/min
Molarity/minute
etc.
As an example, let's look at this reaction:
A + B C
In order to produce any of substance C, molecules A and B must run into each other
(collide). The collision must also have enough energy to exceed the activation energy.
Finally, A and B must be properly oriented when they collide to get the right reaction.
(see figure 4.13 page 116). In summary, to get a reaction you need:
1. collisions
2. collisions with enough energy
3. collisions with the right orientation
The number of collisions per unit time (like per minute) can be increased in several ways.
1. Increase the concentration of one or more reactants.
Example: If 1 mole of A and 1 mole of B are in a closed container, they
can react to make C.
If 2 moles of A and 1 mole of B are put in the same closed
container, the reaction will be faster because the B's can
run into the A's more often.
2. Increase the temperature.
This actually has two effects. First, it makes all the molecules move
around faster, so they simply run into each other more often. The other
effect, is that since they are moving faster, the molecules have more
energy and there is more chance that each collision has the energy needed
to exceed the activation energy.
3. Add a catalyst. A catalyst is a chemical that makes a reaction speed up, but it
is left unchanged by the reaction. They come out of the reaction just like
they went in. They don't get used up.
Most reactions in living systems need catalysts. These are called enzymes
and most of them are proteins.
The main effect of all catalysts (whether biological or not) is to lower the
activation energy of a reaction. This makes is energetically easier for the
reactants to get over the "hill" to change into the products.
An example is a hydrogenation reaction. These are on page 351 in your text. They
require a metal like nickel, platinum or palladium. Here, the metal provides a surface for
the reactants to meet. Many other catalysts for many kinds of reactions are known.
Reversible Reactions
Some reactions can go either way. By this I mean that reactants can be changed into the
products and these products can also be changed back into the original reactants.
Example: If both of these reactions are possible:
A + B C
C A + B
we can write these in a combined form:
A + B ÐC
The reaction of A and B to make C is referred to as the "forward" reaction and the
reaction of C to make A and B is referred to as the "back" or "reverse" reaction.
The symbol: Ð means that there is a dynamic equilibrium. This means that both
reactions continue to occur simultaneously, but that the concentrations of all chemicals
stay unchanged.
Keep in mind:
The word dynamic refers to something changing.
The word equilibirum refers to something not changing.
Here's a "real life" example of a dynamic equilibrium There are 2 rooms (A and B) with
a revolving door in between. We start with some people in room A and a different
number of people in room B. Now let 1 person go from room A to room B per minute.
room A  room B
If nothing else happens then eventually everyone ends up in room B. However, let's also
let people go from room B to room A with 1 person per minute.
room B  room A
Since two processes are happening simultaneously, we'll combine this into a single
"reaction":
room A Ð room B
As long as the "forward" and "reverse" processes occur at the same rate, we'll have a
dynamic equilibrium, no matter how many people we start out with in each room. Let's
say we start with 10 people in room A and 6 in room B. Now let three minutes pass with
people going in each direction at the rate of 1 person per minute. At the end of the three
minutes there will be different people in each room (that's the dynamic part), but there
will still be 10 people in room A and 6 in room B (that's the equilibrium part). We
might call the number of people per room the "concentration of people" and that stays the
same in each room no matter who the people are in each room.
As a chemical example:
carboxylic acid + alcohol Ð ester + water
If you mix a carboxylic acid and an alcohol (and add a little bit of acid catalyst) you can
make an ester and some water. However, in the same container, if the ester molecules
run into water they can reproduce the carboxylic acid and alcohol. If the rates of the
forward and reverse reactions are the same, we'll have a dynamic equilibrium.
Le Chatelier's Principle
This principle states that if you start out with a system that's in a dynamic equilibrium
and then you do something to "disturb" it, the equilibrium will respond so as to mimimize
the disturbance. The system will also re-establish an equilibrium, but a different one that
it started out at.
Next time we'll look at some examples, but here are a few ways we'll "disturb" a system
at equilibrium:
- add a product or reactant to the mix
- take away a product or reactant from the mix
- cool or heat
- for gases we can also change the pressure