1. The Mole and Stoichiometry
Mr. Kinton
Enloe High School
Honors Chemistry

2. Units of Measurement and the Metric System
Scientific work makes observations on qualitative and quantitative data
Some quantities that chemists measure include:
Density
Mass
Volume
Energy

3. Units of Measurement and the Metric System
To help communicate more clearly across the world SI units were introduced
Preferred metric units for use in science
All conversions are done by either multiplying or dividing by some multiple of 10
Prefixes are used to indicate the size of the unit
Without units and a common system, it would be more difficult for scientists to communicate their findings

4. The Metric System

5. Common Units in Chemistry
Grams (g) and kilograms (kg)
Meter (m) and nanometer (nm)
Joule (J), calorie (cal), and electron volt (eV)
Seconds (s) and minutes (min)
Liter (L), milliliter (mL), and microliter (μL)
Mole (mol)
Kelvin (K) and Celsius (oC)
Pascals (Pa), atmospheres (atm), millimeters of mercury (mmHg)
What quantities are being measured with these units?

6. Unit Conversion
Given the following quantities convert them to the desired unit:
22.5 mg to g
15.0 mL to μL
10.5 g to kg
15 GB to MB

7. Scientific Notation
A way to represent large or small numbers using less digits
Example:
4,000 = 4 × 103
= 5 × 10-6
6,325,000 = × 106

8. Recording Measurements and Significant Digits
Chemistry is a laboratory science and data is collected based on laboratory experiments:
How does a Chemist know when they have included enough decimals in their measurement?
Significant figures: the digits that indicate the precision with which a measurement is made
All digits of a measured quantity are significant but the last digit is uncertain

9. Significant Figures We will examine several measurements:
1.350 miles
39.10 grams
810. kJ
In order to determine how many significant figures are included in your measurement check the interval on your measuring device
Lets examine 2 problems in your packet (numbers 1 and 3)

10. Counting Significant Digits in a Measurement
When in lab we have to look at the instrument we are using but in calculations we will need to examine the measured numbers given in our problems
Here are the rules that must be followed when determining how many significant figures a measurement contains:
All numbers between 1-9 are significant
If you have zeros then follow these guidelines

11. Determining Which Zeros are Significant
Zeros between nonzero digits are always significant
Zeros at the beginning of a number are never significant; they are indicating the position of the decimal point
Zeros that are at the end of a number AND come after the decimal point are always significant
When a number ends in zeros but contains no decimal point, the zeros may or may not be significant

12. Determining the Number of Significant Figures
Determine how many significant figures each of the following measurements contains:
87009 km m
85.00 g
7, cm
7.000 × 103 cm

13. Calculations and Significant Figures
When doing calculations you can only be as precise as your least precise measurement
Instead of just choosing a location to round off to at random we use the significant figures of our measurements guide our rounding
Additionally, when recording measurements, your final answer should only contain one uncertain digit

14. Multiplication and Division with Sig Figs
Answers with multiplication and division must be reported with the same number of significant figures as the measurement with the fewest significant figures
6.221 cm × 5.2 cm = cm2 rounds to 32 cm2
23.0 cm × 432 cm × 19 cm = 188,784 cm3 rounds to 190,000 cm3
In rounding your number, look at the leftmost digit to be dropped:
If the leftmost digit to be removed is less than 5 the preceding number is unchanged
If the leftmost number to be removed is 5 or greater, increase the preceding number by 1

15. Addition and Subtraction with Sig Figs
The result can have no more decimal places than the measurement with the fewest number of decimal places
20.4 cm cm + 83 cm = cm or 105 cm
mL mL mL = mL or mL
Rounding rules are the same as with multiplication and division

16. Dimensional Analysis/Factor Label Method
Method of problem solving in which the units are carried through all calculations
Ensures that the final answer has the correct units
Gives a systematic way to solve problems and to help identify possible errors
Conversion factor: a fraction whose numerator and denominator are the same quantity expressed in different units

17. Dimensional Analysis/FLM
If an object is 8.50 in long, how many cm long is the object?
If a woman has a mass of 115 lbs, what is her mass in grams?

18. The Mole and Particles
The amount of matter that contains as many particles as the number of atoms in exactly 12g of Carbon-12
Avogadro’s number: the number of atoms present in a 12g sample of Carbon-12 which is 6.02×1023
1 mol Carbon-12 atoms = 6.02×1023 carbon-12 atoms
1 mol H2O molecules = 6.02×1023 water molecules
1 mol NO3- ions = 6.02×1023 nitrate ions
Using the mole allows us to move from the atomic scale to the macroscopic scale

19. The Mole and Particles
How many moles of magnesium are in 3.01×1022 atoms of magnesium?
How many molecules are there in 4.00 moles of glucose, C6H12O6?

20. The Mole and Molar Mass
A mole will always correspond to 6.02×1023 , however the mass of a mole of 2 different objects will be different
This allows us to make the connection that 1 atom of Carbon-12 has a mass of 12 amu
At the laboratory level, 1 mol of Carbon-12 has a mass of 12g

21. The Mole and Molar Mass
This relationship also works when we examine the formula weights of compounds as well
1 molecule H2O has a mass of amu
1 mol H2O has a mass of 18.02g
Mass in grams of 1 mol of a substance is the molar mass

22. The Mole and Molar Mass
Give the molar mass for the following compounds:
KCl
Ca(NO3)2
CuSO4 ● 5H2O
How many moles are in 28g of CO2?
What is the mass of 5.0 mol of Fe2O3?

23. The Mole and Volume
We can also use the mole to help us determine the volume of a gas at standard temperature and pressure (0o C and 1 atm)
1 mol of any gas = 22.4 L
With the mole we can convert between particles, mass, and volume.

24. The Mole and Volume
Determine the volume, in liters, occupied by 2.0 moles of a gas at STP.
How many moles of argon are present in 11.2 L of argon gas at STP?

25. Conversions between Moles
While it is useful to convert to moles sometimes we will want to move to another unit
Follow this process: Given unit  Moles  Desired Unit
How many glucose molecules are in 5.23g of C6H12O6?
How many O atoms are in a sample of 4.20g of HNO3?

26. Percent Composition
When given a formula we know the number of atoms or ions present in that compound
Percent composition by mass will allow us to determine the empirical formula of a compound
Empirical formula: compound which shows the relative number of atoms in a compound

27. Percent Composition
There are 2 ways to determine the percent composition of a compound:
% Composition = (subscript of element × molar mass of element) ÷ molar mass of compound
% Composition = mass of component in a sample ÷ total mass of the sample
Anytime you determine the formula of a compound using percent composition it will be an empirical formula

28. Percent Composition
Ascorbic acid contains 40.92% C, 4.58% H, and 54.50% O by mass. What is the empirical formula of ascorbic acid?
A 5.325g sample of methyl benzoate, a compound used in the manufacture of perfumes is found to contain 3.758g of carbon, 0.316g of hydrogen, and g of oxygen. What is the empirical formula?

29. Molecular Formula
Tells us the actual number of atoms of each element in the compound
Will always be a whole number multiple of the empirical formula
In order to determine the molecular formula, you must know the molecular weight of your compound
This is done by comparing the empirical formula weight to the molecular formula weight

30. Molecular Formula Step 1: Find the empirical formula
Step 2: Calculate the formula weight of the empirical formula
Step 3: Find the molecular weight given in the problem
Step 4: Divide the molecular weight by the formula weight. Round to a whole number
Step 5: Multiply all subscripts in the empirical formula by the whole number

31. Molecular Formula
Earlier we determined the empirical formula of ascorbic acid to be C3H4O3. The experimentally determined mass of this compound is 176 amu. Find the molecular formula.
Ethylene glycol, the substance used in antifreeze, is composed of 38.7% C, 9.7% H, and 51.6% O by mass. Its molar mass is 62.1 g/mol. What is the molecular formula for ethylene glycol?

32. Stoichiometry
Quantitative study of reactants and products in a chemical reaction
As a result, you must have a balanced chemical equation in order to perform a stoichiometry calculation
The coefficients in the equation tell you how many moles of each substance must combine to form products
In the reaction: 2H2 (g) + O2 (g)  2H2O (l)
2 moles of hydrogen reacts with 1 mole of oxygen to form 2 mole of liquid water
If we can balance an equation, we can determine a myriad of chemical quantities based on that equation.

33. Stoichiometry
As we solve problems using stoichiometry follow these steps:
Write the correct formulas for all reactants and products. Balance the equation
Convert the quantities of the given or known substances into moles
Use the mole ratio (coefficients) from the balanced equation to convert to the moles of the substance needed
Convert the moles needed to the unit that is specified in the problem
Check that you answer is reasonable

34. Stoichiometry
How many grams of water are produced in the oxidation of 1.00 g of glucose?
Calculate the volume of CO2 produced by the combustion of 1.00 g of butane.

35. Stoichiometry and Limiting Reactants
During most reactions a reaction will stop once a reactant is completely used up
Limiting reactant/reagent: the reactant that is completely consumed during the course of a chemical reaction
Excess reactant/reagent: reactant that is present in a greater quantity and is left over once the reaction stops

36. Limiting Reactant
Urea is prepared by reacting ammonia and carbon dioxide. Water is the other product. Consider the reaction of g of NH3 with 1,142 g of CO2.
What is the maximum mass of urea that can be formed?
What is the limiting reactant?
How much of the excess reactant remains?

37. Limiting Reactant
Consider the reaction between aluminum and iron (III) oxide. In one process, 124 g of aluminum reacts with 601 g of iron (III) oxide.
What is the maximum mass of iron that can be produced?
What is the limiting reactant?
How much of the excess reactant will remain once the reaction has finished?

38. Reaction Yield
Anytime a reaction takes place it will yield a certain amount that is less than what we would expect
This is due to side reactions that take place and our inability to capture all products
Reaction yield or percentage yield is the amount of product expressed as a percentage of the actual yield and the theoretical yield of the products

39. Reaction Yield
Actual Yield: the amount of product that is experimentally produced
Theoretical Yield: the calculated amount from the balanced equation.
This will be the amount of product formed based on the limiting reactant
% Yield = (Actual yield ÷ Theoretical yield) × 100

40. Reaction Yield
In a certain industrial operation, 3.54 × 107 g of TiCl4 reacted with 1.13 × 107 g of magnesium.
Calculate the theoretical yield of Ti
Calculate the percent yield of Ti if 7.91 × 106 g were obtained in the operation

41. Reaction Yield Consider the reaction of vanadium (V) oxide.
Calculate the theoretical yield of vanadium in a process that involves the reaction of 1.54 × 103 g of vanadium (V) oxide with 1.96 × 103 g of calcium.
What is the percent yield if 803 g of vanadium were obtained in the process?