###### Chem 2045   Return  Chapter 3 Stoichiometry: Calculations with Chemical Formulas and Equations

Chemical equations:
1. Formation of ozone,

NO2 --> NO(g) + O(g)

O(g) + O2(g) --> O3(g)

Be familiar with terms such as reactant and product. Be able to interpret and draw, on a molecular level, what is being "said" by the chemical reaction. For example, the reaction of an oxygen atom an an oxygen molecule to give ozone can be indicated as:

2. Combustion of hydrogen. This reaction is the reverse of the electrolysis of water that I showed you using the battery and salt water. Do you think you know what "combustion" means? Try writing down a definition, then look it up in the text to see how you did.

2H2(g) + O2(g) --> 2H2O(l)

Be able to balance simple chemical equations. Remember, coefficients (in front of the compound) are used for balancing. Never change the subscripts within a formula when balancing an equation.

Try this: balancing the following reactions
1.
The combustion of ethane. Note that ethane, C2H6(g) only has hydrogen and carbon in it. Such compounds are called hydrocarbons. The complete combustion of any hydrocarbon gives only CO2(g) and H2O(l) as its products. This process is similar to the way our bodies get fuel. The food we eats reacts with O2 (that's why we inhale) to form CO2 (that's why we exhale) and H2O.

C2H6(g) + O2(g) --> CO2(g) + H2O(l)

Sometimes this reaction is balanced using fractions. For now we will use whole numbers.

2. The reaction of calcium metal with water.

Ca(s) + H2O(l) --> Ca(OH)2(aq) + H2(g)

Of course, we don't write the coefficient "1". We just leave it blank and we assume it is one. (see the equations above for formation of ozone.) Compare this reaction with equation 3.6 in you text book. How is it similar? Different? Why does the calcium metal end up with with 2x OH's? Can you detect a trend? Try predicting the reaction of barium, Ba, with water. Check out a movie of the reaction of sodium metal with water. See how useful the periodic table can be at suggesting patterns of chemical reactivity. Scientists rely on these patterns to help make sense of the infinite number of chemical reactions.

3. The reaction of melted aluminum with solid barium oxide.

Al(l) + BaO(s) --> Al2O3(aq) + Ba(l)

You should also be able to identify combination and decomposition reactions.
4. For the following reaction, write and balance the chemical equation and then identify whether it is a combination or decomposition reaction.

The reaction when potassium hydrogen carbonate is heated to give solid potassium oxide, water and carbon dioxide.

Atomic and molecular weights:
Obviously, even a tiny sample of matter ( maybe 1.0 g) has countless molecules, atoms or ions in it. How are we supposed to know how many? Well, there are several ways to indirectly determine how many there are of a very large number of small objects. For instance, check out this site at Nature. The problem is that we can't count out molecules from a jar. For instance, if I want to run the reaction in 3. above, how much BaO do I need if I have 54.0g of Al? The solution to this problem lies is setting an arbitrary standard. The chosen standard is carbon-12. It is agreed that 1 atom of the isotope of C-12 is defined as exactly 12 amu.

Mass spectroscopy:
This technique is used to find the relative masses (compared to C-12) and % abundance of the isotopes of a given element or to find the relative masses of any other element on the periodic table. Once we know how much of there is of a given isotope, for example of Cu, we can take the % weighted average and determine the atomic mass of that element. This is the value that is reported on the periodic table.

Try this: calculate the atomic weight of copper, given that a natural sample typically consists of 69.17% copper-63, which has an atomic mass of 62.94 amu, and 30.83% coper-65 which has an atomic mass of 64.93 amu.

The atomic masses compiled on the periodic table can be used to determine the formula weight of a compound and the relative amounts, by mass, of each of the elements that makes it up.

What is the formula weight of Al2O3?

What is the % composition of each element in Al2O3?

The mole:
We defined the mole (a.k.a Avogadro's number, N) to equal the # of carbon atons in exactly 12 g of pure C-12. Today we know that a mole of C-12 has about 6.0221421 x 10^23 atoms. Once we know how many things are in a mole, we can have a mole of anything. For example:

1 mole of cars = 6.02 x 10^23 cars
1 mole of students = 6.02 x 10^23 students
1 mole of Na atoms = 6.02 x 10^23 atoms
1 mole of H2O molecules = 6.02 x 10^23 molecules
1 mole of K+ ions = 6.02 x 10^23 ions

Obviously a mole is a huge number and is only useful for very small things that we might have a lot of. In the same way that a dozen = 12 things and is useful for counting eggs or bagels, but wouldn't be useful for counting atoms because a dozen is too small. The mole leads to the concept of molar mass which is the mass of 1 mol of an object (ie. 6.02 x 10^23 of that object). The molar mass of an element (in g/mol) is numerically equal to the atomic weight of that element (in amu). For example:

C has a molar mass = 12.01 g/mol
Na has a molar mass = 22.99 g/mol
H2O has a molar mass = 18.02 g/mol
OH- ion has a molar mass = 17.01 g/mol

1.
What is the molar mass of carbonic acid, H2CO3?

2. How many grams would you need to have 2.50 mol of H2CO3?

3. How many molecules of H2CO3 are there in the 2.50 moles?

4. How many O atoms are there in 2.50 moles of H2CO3?

5. The neutralization reaction for KOH by H2CO3 is as follows: KOH (aq) + H2CO3 (aq) -> H2O + K2CO3(aq)
How much KOH is required to neutralize 155 g H
2CO3?

Molar mass:
On Friday, we defined the molar mass as the mass of 1 mol of an object (ie. 6.02 x 10^23 of that object). The molar mass of an element (in g/mol) is numerically equal to the atomic weight of that element (in amu). This makes things really easy since we can get the stomic weight right off the periodic table.

The most powerful application of this concept is that it allows us to measure quantities of atoms without having to count them. For example, let's say I want to perform the following reaction for the production of ammonia which is used primarily in fertilizers and is one of the top ten chemicals produced in the United States (over 16 billion kg were produced in 1995!):

N2(g) + H2(g) -> NH3(g)

It should be clear from the balanced reaction that I need 3 times as much H2 as N2. Since the N2 and H2 don't weigh the same, we have had no way of knowing how much of each gas was needed to give the correct ratio. NOW, using the molar mass of N2 and H2, we can determine the correct ratio that we are required to weigh out. As an example, imagine we have 9.50 g of N2(g), how much H2 do we need?

Limiting reactants:
You will find that in chemistry, you don't always have exactly the right amounts of each chemical that are required for the stoichiometry of the reaction. For example, with the above reaction, we saw that for 9.50 g of N
2, we should have 2.05 g of H2. If I tried to run the reaction with 9.50 g of N2, but only 1.00 g of H2, I would run out of H2 before all the N2 was used up. H2 is called the limiting reactant.

Try the following:
1. Methanol, CH
3OH is an excellant fuel and can be made by the reaction of carbon monoxide and hydrogen

CO (g) + H2(g) -> CH3OH (l)

Suppose 356 g of CO is mixed with 65.0 g of H2.

a. Which is the limiting reactant?

b. What is the maximum mass of methanol that can be formed?

Empirical formula from analysis:
Scientists are often faced with identifying unknown compounds or determining the formula for a newly made compound. A new drug for instance might be sent off to a lab to be analyzed. Typically, the results of such an alaysis are the percentage by mass of each element in the unknown. It is then possible to determine the empirical formula from analysis.

Try the following example:
1. The elemental analysis of an unknown ionic compound gave the following results: 18.8% Na, 29.0 % Cl and 52.2%O. What is the empirical formula and name of the compound?

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