Stoichiometry Atomic Mass (atomic weight)
• Atoms are so small, it is difficult to discuss how much they weigh in grams
• We use atomic mass units – an atomic mass unit (AMU) is one twelfth the mass of the catbon-12 atom (the standard)
• This standard gives us a basis for comparison
• The decimal numbers on the periodic table are in amu’s Example 1
What is the atomic mass for carbon? For zinc?
They are not whole numbers
• because they are based on the average mass of the common isotopes of an atom
• To calculate, multiply the % abundance of each isotope to the relative mass of each and add together
Example 2 & 3
There are 2 isotopes of bromine, one with a mass of 78.918336 amu and an abundance of 50.69%. The other isotope has a mass of 80.916289 amu and an abundance of 49.31%. Calculate the atomic weight (average atomic mass) of bromine.
There are 2 isotopes of nitrogen, one with an atomic mass of 14.0031 amu and one with a mass of 15.0001 amu. What is the percent abundance of each?
The Mole (mol)
• A large number used for counting atoms
• It is the number of particles in exactly 12 grams of carbon 12.
• Equal to 6.022 x 1023 of anything, no matter what the substance
• 6.022 x 1023 is called Avogadro’s number
Example 4-5
Convert 1.35 moles of carbon disulfide(CS2) into molecules?
How many moles are in 3.59 x 1021 formula units of sodium nitrate, NaNO3
Molar Mass (molecular weight)
• Mass of 1 mole of a substance
• For elements, it is the amount in grams numerically equal to the atomic mass in amu’s
• For compounds, add up the individual atomic masses of the elements that make it up. Remember a subscript in a chemical formulas tells you how many atoms of an element you have and this must be accounted for
• Represented by M and has the units grams/mol Example 6-9
Calculate the mass of a carbon atom.
Calculate the mass of a zinc atom.
**** Notice that the answers for the questions above are the same as the answers to question 1 except the units have changed.****
What is the formula mass of chloric acid?
Calculate the molar mass of Al2O3 in grams/mol
1 mol = 6.022 x 1023 particles (atoms, molecules, formula units)
1 mol = molar mass (g)
Conversions using Molar Mass Examples 10-13
How many moles are there in 1.00 g of Al2O3?
Determine the number of moles of calcium carbonate in a stick of chalk weighing 14.8 g.
What is the mass in grams of .287 mol of acetylsalicylic acid found in aspirin, C9H8O4.
Calculate the atoms in 11 grams of C2H6
Percent Composition
• Percent that each element in a compound is composed of
• Calculate the molar mass of the compound. Then divide the total mass of each element, by the molar mass(total mass of the compound), and multiply by 100.
• Check you answer to see if the total of the percent masses of all elements equals 100%
Example 14-15
Calculate the mass percent of each element in Al2O3.
Calculate the percent composition of water in the hydrate, CaSO4 · 2H2O
Empirical Formula
• The lowest ratio of atoms in a molecule
• From % composition, you can determine the empirical formula, based on mole ratios
• Remember this rhyme: Percent to mass, mass to mole, divide by small, multiply till whole (last step is not always mandatory)
• Steps to calculating empirical formula
1. If you are given a % composition, assume you have 100 grams and switch the
% unit to g. If you are given the actual number of grams of each element, use those instead
2. Convert the grams of each element to moles of each element.
3. Divide each mole amount by the smallest mole amount.
4. If your answers are within .1 of whole number, round up or down and these numbers become your subscripts.
5. If your answers are not within .1 of a whole number, you must find the smallest number that you can multiply by all the mol amounts you have just calculated to get them to be all whole numbers or within .1 of a whole number.
Example 16 -18
A sample of compound is made up of 78.20 g of potassium and 32.06 g of sulfur, what is the simplest formula (empirical)?
A 25.00 g sample of an orange compound contains 6.64 g of potassium, 8.84 g of chromium and 9.52 g of oxygen. Find the empirical formula.
A compound has the following percent composition: 31.9%K, 29.0 % Cl and 39.1%O. What is the empirical formula?
Molecular Formula
• Empirical means the lowest ratio, molecular means the actual ratio of the molecule that exists.
• To calculate, you need the molar mass of the molecular formula and the empirical formula’s mass(you will probably have to calculate this
• Steps to calculate the molecular formula:
1. Divide the Molecular formula mass by the empirical formula mass to get a whole number.
2. Multiply the whole number(factor) to the subscripts of the empirical formula Example 19
Propene has the simplest formula CH2. What is the molecular formula if it has a molar mass of 42.09 g/mol?
Chemical Equations
• Sentences that describe what happens in a chemical reaction
• Reactants are the starting substances found on the left of an equation
• Products are the new substances created which are found on the right of an equation
• Reactants Products (yields)
• Law of Conservation of Mass: Matter can neither be created nor destroyed, just changed in form.
• Equations should be balanced in order to satisfy the law of conservation of mass- same number of each kind of atom on both sides of the equation must exist
__ CH4 + __ O2 __ CO2 + __ H2O
1 C 1 4 H 2 2 O 3
Important Abbreviations : (s) solid (l) liquid (g) gas (aq) aqueous heat ∆ catalyst
• Hints for balancing:
1. Never start with hydrogen or oxygen
2. Treat polyatomic ions as chunks not separate atoms
3. Do not put a coefficient of 1 down until the end when all atoms are balanced 4. If you have an odd number of atoms on one side and an even number of those
atoms of the other, place a 2 as a coefficient next to the chemical formula with the odd numbers to make it even.
Examples 20-23
__ Ca(OH) 2 + __ H3PO4 __ H2O + __ Ca3(PO4) 2
__ Cr + __ S8 __ Cr2S 3
__ KClO 3 ___ KCl + __ O 2
Translate this and then balance: Solid iron(III) sulfide reacts with gaseous hydrogen chloride to form solid iron(III) chloride and hydrogen sulfide gas
Meaning of the coefficients in a Chemical Equation
• Used to describe a reaction in moles, and particles (molecules, formula units and atoms) but not grams. For example: 2 H2 + O2 2 H2O would be interpreted as 2 molecules of hydrogen react with 1 molecule of water to produce 2 molecules of water OR 2 moles of hydrogen react with 1 mole of oxygen to produce 2 moles of water!
• A Mole Ratio is the conversion factor between 2 different amounts in a balanced chemical equation
2 KClO 3 2 KCl + 3 O 2 Possible Mole Ratios could be:
• Given an amount of either reactant or product, you can determine the other quantities in a reaction using dimensional analysis
• Use conversion factors from molar mass ( ? grams = 1 mole) and the balanced chemical equation mole ratio (? Mol A = ? Mol B)
• 4 types of problems
1. mole –mole: 1 step process using a mole ratio to convert from moles of substance A to moles of substance B
2. mole-mass: 2 step process using a mole ratio to convert from moles of substance A to moles of substance B and then the molar mass conversion factor to convert between moles of B to grams of B 3. mass-mole: 2 step process using the molar mass conversion factor to convert from mass of A to moles of A and then the mole ratio to convert from moles A to moles of B
4. mass-mass: 3 step process using the molar mass conversion factor to convert from mass of A to moles of A, then a mole ratio used to convert from moles A to moles of B, then a molar mass conversion factor to convert between moles of B to mass of B
2 mol KClO 3 = 2 mol KCl OR 2 mol KCl = 3 mol O 2 OR 2 KClO 3 = 3 mol O 2
Example 24-25a-c
(Mole To Mol) How many moles of sulfur trioxide can be produced if 8.1 moles of oxygen react with sulfur. 2S + 3 O2 2 SO3
(Mass To Mole) One way of producing oxygen O2 involves the
decomposition of potassium chlorate into potassium chloride and oxygen gas.
A 25.5 g sample of potassium chlorate is decomposed. How many moles of O2 are produced?
2 KClO 3 2 KCl + 3 O 2
B. (Mass to Mass) How many grams of O2?
C. (Mass to Mass)How many grams of potassium chloride?
Limiting Reagent
• Limiting reagent is the reactant that determines the amount of product formed ( it is the one you run out of first)
• Excess Reagent The reactant that you don’t run out of.
• How to tell if the problem you are working is a limiting reagent: IF 2 AMOUNTS OF REACTANT ARE GIVEN TO YOU IN THE PROBLEM!
• Steps to determine the LR(limiting reagent)
1. Take each reactant amount and convert to moles. If already in moles move to the next step.
2. Divide each mole amount by the coefficient of that substance from the balanced chemical equation.
3. Whichever is the smallest number will identify the limiting reagent (LR)
4. Use the original amount of the reactant from the problem to begin your conversion process to determine either how much excess reagent you have left over or how much product you will form.
Example 26
Ammonia is produced by the following reaction: N2 + 3 H2 2 NH3
What mass of ammonia can be produced from a mixture of 100. g N2 and 500.
grams of H2?How much excess reagent do you have left over?
Percent Yield
• Percent Yield is the ratio in percent of amount of product you produced in the lab compared to the amount of product that you should of produced if the reaction went to completion and no problems arose in the lab
• Must always use the LR to determine the amount of product produced
• Theoretical yield- is the amount you would produce if everything went perfect;
use the balanced chemical equation for this one
• Actual Yield is what you make in the lab under imperfect conditions
• % yield = Actual/Theoretical x 100% OR
• % yield = what you got/what you should of got x 100%
Example 27
Aluminum burns in bromine producing aluminum bromide. In a lab, 6.0 g of aluminum reacts with excess bromine. 50.3 grams of aluminum bromide are produced. What are the three types of yield? 2 Al + 3 Br2 2 AlBr3
Special type of Conversion: Combustion Analysis
Steps to Solving a problem involving C, H and O in the formula :
1. Convert the mass of carbon dioxide to moles of carbon dioxide . Then convert to moles of carbon. Then convert to grams of carbon
2. Convert the mass of water to moles of water. Then convert to moles of hydrogen. Then convert to grams of hydrogen.
3. Subtract the mass of C and H from the mass of the sample to get the mass of O 4. Divide all the mole amounts by the smallest number to get the subscripts.
5. If the amounts are not within .1 of a whole number, you will have to multiply till whole like when doing the typical empirical problem.
Example 28
The compound that gives fermented grape juice, malt liquor, and vodka their intoxicating properties is ethyl alcohol, which contains the elements carbon, hydrogen, and oxygen. When a sample of ethyl alcohol is burned in air it is found that 5.00 g ethyl alcohol 9.55 g CO2 + 5.87 g H2O. What is the empirical formula?
