CH 12 NOTES
Review
The macroscopic vs. the submicroscopic
Is it practical to count out 200 grains of rice, or 200 pieces of macaroni? NO WAY! We sell these items by the pound. How would we be able to calculate how many items are in a container without counting?
In your group, design a method to figure out how many items are in each bag, without counting. What information would you need? Write one complete description of your method per group.
DEMO: Counting by Weighing
1.How many sourballs are in the jar?
What you need; the data:
Mass of the empty jar (with stopper) = 209.624 g
Mass of the full jar (with stopper) = 396.227 g
Average mass of one sour ball = 3.92 g
2.How many atoms are in 1.0079 g of Hydrogen if the average mass of one atom is 1.67425 x 10-24 g?
Section12.1
STOICHIOMETRY : the study of the relationships between the macroscopic quantities like mass and volume, and the submicroscopic quantities like number and mass of the atoms involved in a chemical reaction.
What is a mole?
The group or unit of measure that we use in chemistry to count numbers of atoms, ions, molecules or formulas units of a substance (elements or compounds)
What is Avogadro’s Number? 6.02 x 1023
This HUGE quantity represents the number of particles in one mole of any substance. A mole can be analogous to a dozen, a pair, a century…you know how many individual parts there are in each one of these measurements, right? 12, 2, 100…a mole is just A LOT more.
DEMO: What is the same about moles?
What is Molar Mass?
The mass in grams of one mole of an element or a compound!
For elements, it is the atomic mass in units of grams
Ex. One atom of magnesium = 24.31 u (atomic mass units)
One mole of magnesium = 24.31 g
For compounds, it is the sum of all the masses of each element in the compound’s formula
Ex. One mole MgSO4 = 24.31 + 32.07 + 4(16.0) = 120.4 g
So what about molecular mass and formula mass?
These are not hard to calculate, because essentially the value of each is the same. It is the units that are different!
Ex. Formula Mass is a term used for IONIC compounds
NaCl has a FORMULAmass of 58.45 u and a MOLAR mass of 58.45 g ~ Same quantity, different unit
Ex. Molecular Mass is a term used for COVALENT compounds
C2H6 has a MOLECULAR mass of 30.1 u and a MOLAR mass of 30.1 g
When solving problems like the ones below, you must use DIMENSIONAL ANALYSIS (Factor Label Method).
Practice Problems, p. 410
#s 1, 2c, 2d, 3c, 3d, 4c and 4d
Extension Problem: If the circumference of the Earth is 4.0 x 109 cm, and a paper clip is 3.0 cm long, how many times would 1 mole of paper clips go around the Earth?
Section 12.2 Using Moles
Objective: Predict quantities of reactants and products in a chemical reaction
DEMO: Vegetable Soup Problem
The following recipe serves 4:
4 potatoes (.3 lb each)
2 onions (.2 lb each)
8 carrots (.1 lb each)
4 stalks of celery (.05 lb each)
1.4 lbs of water
What is the total mass of the soup?
How would you make enough for 8?
How about 1240?
DEMO: Mole Trail Mix
RESULT:
No matter how we change the quantity, the proportion of ingredients always stays the same! This is also true in chemical reactions.
**With a balanced chemical equation and number of moles, we can predict the exact amount of reactant and product in a reaction**
There are 4 steps to follow:
- Write the balanced chemical equation
- Convert the given mass or volume to moles
- Use the coefficients in the chemical equation to set up a mole ratio (the coefficients are the # of moles!)
a. HINT: The substance you are solving for goes on TOP
- Convert these moles back to mass or volume as required
Ex: N2 + 3H2 2NH3
This reaction can be stated as: 1 mole of nitrogen will react with 3 moles of hydrogen to produce 2 moles of ammonia.
What if we had 1 mole of nitrogen and only 1.5 moles of hydrogen? Which is the limiting reactant?
Practice Problems, p. 415
#10: predicting the mass of a product
#11 and 12: predicting the mass of a reactant
MOLAR VOLUME (gas): the volume that one mole of gas occupies at STP; 22.4 L
Practice Problems, p. 416
#13 and #14: Find the mass/volume of reactant/product from a given quantity (Both are at NON-STANDARD conditions!)
IDEAL GAS LAW : relates pressure (P), volume (V),
temperature(T) and # of particles or # of moles (n) of a gas; Use the formula:
PV = nRT
R is a constant = 8.31 kPa . L
mol . K
** If the given pressure is in kPa, use the value for R above. If the given pressure is in atm, then use the value
R = .08205 atm . L
mol . K
** If the given pressure is in mm Hg, then use the value
R = 62.36 mm Hg . L
mol . K
** YOU WILL NOT HAVE TO MEMORIZE THESE!**
Practice Problems, p. 419
#s 15-17
Do NOT let the algebraic formula get you! These are all “plug & chug” problems.
THEORETICAL YIELD AND ACTUAL YIELD
In an experiment, the theoretical yield is pre-determined by your stoichiometric calculations. The actual yield is what you really get when the experiment is over.
The reasons that these numbers are different from one another are: poor collection techniques, faulty apparatus, lack of time, and just plain old human error. (Think of how much salt you lost in your reaction in the lab!)
PERCENT YIELD
This value is used to express the efficiency of your reaction, i.e. how close were you?
EX: If you calculate that you should yield 2.50 g NaCl in your experiment, but you actually get 2.10 g, you can calculate your percent yield by using this formula: actual yield x 100%
theoretical yield
Example: 2.10g/2.50g x 100% = 84.0%
**This means that you obtained only 84.0% of the possible product, and you may have committed an error somewhere.
Objectives:
- Determine mass percent of an element in a compound
- Identify formulas of compounds by using mass percents and mole/mass ratios.
DEMO/HANDS-ON ACTIVITY: EVERDAY MASS PERCENTAGE
Problem: What percentage of the gum you chew is sugar?
Design an experiment that would physically allow you to determine this amount.
Analysis:
- What is the percentage of sugar?
- What is the molar mass of sugar, C12H22O11?
- Determine the number of moles in the mass of sugar your group consumed.
- How many molecules of sugar are in the mass of sugar your group consumed?
DETERMINING MASS PERCENTS
- Mass percent refers to the relative mass that each element contributes to the total mass of a compound.
For example: Cl2 is 100% chlorine
- To calculate mass percent, use the following:
Mass of each element x 100%
Total mass of the compound
For example, what is the mass percent of NaCl?
% Na = mass of 1 mole sodium x 100% = 22.99 g x 100% = 39.34% mass of 1 mol salt 58.44 g
% Cl = mass of 1 mole chlorine x 100% = 35.45 g x 100% = 60.66% mass of 1 mol salt 58.44 g
**The percentages should always add up to 100% (+/- .2 %)**
Example:
Calculate the mass percentage of oxygen in the compounds sodium nitrite (NaNO2) and hydrogen peroxide (H2O2).
Both of these compounds contain 2 moles of oxygen
atoms per mole of the compound. Why are the mass percentages of oxygen so different?
DETERMINING CHEMICAL FORMULAS
EMPIRICAL FORMULA: the formula of a compound that has the smallest whole-number ratio of atoms. For example, given the molecular formula for glucose, C6H12O6, you can see that itsempirical formula would be CH2O
When solving for an empirical formula, follow these steps:
EMPIRICAL FORMULAS (E.F.)
Are you given the mass ofeach element in the compound? If yes, then Step 1: Convert/Calculate # of moles of each from the mass
Step 2: Find the MOLE RATIO(If the ratio is not in whole#s, make it so by dividing by the smallest number of MOLES.)
Step 3: Write the Empirical Formula using your ratio
Variation~Are you given mass percent of eachelement? If yes, then assume the sample has a mass of 100 g, then calculate the mass of each element, and return to Step 1 above.
MOLECULAR FORMULAS (M.F.)
If you are given the E.F. and the molecular mass, then go to:
Step 1:Calculate the E.F. mass
Step 2:Determine the MASS RATIO : Divide molecular mass
E.F. mass
Step 3: Multiply the E.F. Ratio by the whole # determined in Step 2. (This is your molecular formula ratio.)
Step 4:Write the molecular formula (you may check your work by calculating the molecular mass of your formula).
HYDRATE FORMULAS
If you are given the mass of a compound and that of water, then:
Step 1:Calculate the # of moles for both the compound & water
Step 2:Find the RATIO
Step 3:Write the formula for the hydrate