1. The measurements/math of chemistry
Unit 5: Stoichiometry
The measurements/math of chemistry

2. Chemical reactions A chemical reaction:
One or more substances are made into one or more NEW substances (bonds are broken and/or bonds are made).
AgI (aq) + F2 (g)  AgF (s) + I2 (s)
The arrow is read “reactants yields products.
Aqueous Silver iodide plus fluorine gas yields solid silver fluoride plus solid iodine

3. Types of Reactions (Rxns)
Synthesis Reactions:
Reactants and Products:
General Equation:
(ex)
Decomposition Reactions:
Combustion Reactions:
Reactants and Products (Complete Combustion):
Incomplete combustion:
2 elements or cmpds  1 cmpd
A + B  AB
2H2 + O2  2H2O
1cmpd  2 elements or cmpds
AB  A + B
2NaCl  2Na + Cl2
Organic cmpd + O2  CO2 + H2O
Organic + O2  CO + H2O
C6H12O6 + O2 
CO2 + H2O
C6H12O6 + O2  CO + H2O
**Only identify

4. Types of Reactions (Rxns)
Single Replacement Reactions:
Reactants and Products:
General Equation:
(ex)
**All other reactions happen automatically  single replacement reactions DO NOT ALWAYS REACT.
**They only occur if the element that is by itself is more reactive than the ELEMENT IN THE SAME COLUMN ON TABLE J in the compound.
(ex) NaCl (aq) + Li 
(ex) NaF (aq) + Cl2 
element + ionic cmpd (aq) (or acid)
 element + ionic cmpd
A + BY (aq)  B + AY
AuCl (aq) + Cs  CsCl + Au
No reaction

5. Types of Reactions (Rxns)
Double Replacement Reactions:
Reactants and Products:
General Equation:
(ex)
2 ionic cmpds (aq) (or acid)
 2 ionic cmpds
AX (aq) + BY(aq)  BX + AY
KCl (aq) + NaNO3 (aq)  NaCl + KNO3

6. Predicting Phases
We should know the phase of any element- we labeled them on our Periodic Tables.
H2O and CO2 are both gases in combustion reactions due to the heat produced in the reaction.
Ionic compounds where there is NO WATER present are solids.
Ionic compounds where there IS WATER present, check Table F.

7. Precipitates:
Precipiates: (usually from single or double replacement reactions)
A solid formed as a product of a reaction that does not dissolve in water (insoluble)
“if you’re not part of the solution, you’re part of the precipitate” 
Table ___________
(ex) Cu(NO3)
(ex) PbCl2
(ex) Ca3(PO4)2
(ex) Ba(OH)2 F
(aq)
(s)
(s)
(aq)

8. Predicting Products Practice
Li3(PO4)(aq) + Mg(NO3)2(aq) 
Sr(s) + NaCl(aq) 
C4H12(g) + O2(g) 
Al(s) + HCl(aq) 
CH3OH(g) + O2(g) 
NaOH(aq) + MgSO4(aq) 
K (s) + Br2(l) 
Al2O3 
Determine the type of reaction.
Write the products based on the reaction type.
Determine NEW subscripts for each compound based on charges.
Determine phase of each product.
BALANCE

9. Law of Conservation of Mass
Mass is neither created nor destroyed in a chemical change.
~ 1789 Antoine Lavoisier
This is shown by a balanced reaction, or having the same mass on the reactants side as the products side.
(ex) 2H2O  2H2 + O2
2.54 g g
0.28 g x
x = 2.26 g 13

10. Balancing Reactions
Therefore, the mass of the reactants = the mass of the products (no atoms can disappear or appear out of nowhere)
The total number of atoms of each element must be equal on both sides of the reaction arrow.
**Balance the elements one at a time using the coefficients. After you have written the correct formulas for all compounds –
NEVER, NEVER, NEVER change the subscripts.

11. Balancing Practice (ex) ____Ag + ____AuCl3  ____AgCl + ____Au
(ex) ____Pb(NO3)2 +____NaCl ____PbCl2 + ____NaNO3
(ex) ____CH4 + ____O2  ____CO2 + ____H2O
(ex) ____C8H18 + ____O2  ____CO2 + ____H2O

12. The Mole Calculating the mass of a compound:
** First you have to know how much of the compound you have and what compound it is to calculate the mass.
The mole (measures how much):
Is a word that means a number (6.02 x 1023) JUST LIKE A DOZEN
(Ex) 1 dozen donuts = _______________________________
1 mole of donuts = ______________________________
(Ex) 1 dozen particles = _______________________________
1 mole of particles = ______________________________
Particle can mean 1 molecule, 1 atom, 1 ion.
12 donuts
6.02 x 1023 donuts
12 particles
6.02 x 1023 particles

13. GFM Gram Formula Mass (GFM)
The mass of 1 mole (6.02 x 1023 atoms/molecules) of an element or compound
C atomic mass: (mass of 1 atom) _________________
C molar mass: (mass of 6.02 x 1023 atoms) ______________
Ar atomic mass: _________________
Ar molar mass: __________________
To find the mass of a mole of a compound you must add the masses of all of the elements in the compound or molecule.
(ex) K(NO3)
(ex) CH4
(ex) Fe3(PO4)2
12 amu
12 g
40 amu
40 g
= (16)
= 101 g
= (1)
= 16 g
= 3(56) + 2(31) + 8(16)
= 358 g

14. Moles of Atoms
If there are a dozen cars in the parking lot, how many dozen wheels are there?
If I have a mole of H2O and each water molecule has 2 hydrogen atoms, how many moles of hydrogen do I have?
Subscripts actually tell you the number of moles of each atom.
(Ex) 1 mole of SO3
How many moles of S atoms?
How many moles of O atoms?
How many moles of atoms in 1 mole of SO3?
What is the mass of 1 mole of SO3?
(Ex) 1 mole of Ca(OH)2
How many moles of Ca atoms?
How many moles of O atoms?
How many moles of H atoms?
How many moles of atoms in 1 mole of Ca(OH)2?
GFM of Ca(OH)2 ? 1 1 3 2 2 4 5
= (16)
= 80 g
= (16) + 2(1)
= 74 g

15. Converting moles   grams:
Table T for mole calculations.
(ex) In the lab you are given 96g of NaOH. How many moles of NaOH do you have?
(ex) You are asked to measure out 4.23 mol of K2CO3, how many grams is this?

16. Converting moles   liters:
1 mole of any gas equals 22.4L
(ex) How many liters are in 2.37 mol of H2?
(ex) How many moles are in 65.9 liters of CO2?

17. Converting moles   #particles:
1 mole of any substance contains 6.02 x 1023 particles
(ex) How many molecules of water are in 2.50 moles of water?
(ex) How many moles of SO3 are present if there are 9.03 x 1023 particles of SO3?

18. Mole Map
Grams
Particles
1 mole = GFM
1 Mole
Liters

19. Mole Map Grams Particles 1 Mole Liters 1 mole = GFM
1 mole = 22.4 L of any gas
Liters

20. Mole Map Grams Particles 1 Mole Liters 1 mole = GFM
1 mole = 6.02 x 1023 particles
1 Mole
1 mole = 22.4 L of any gas
Liters

21. Avogadro’s Hypothesis
Any 2 samples of gas at the same temp, pressure and volume have the same number of particles
(ex) Which of the following would occupy the same volume as 2.5 mol of O2 at STP?
(a) 1.5 mol of CO2 at STP
(b) 3.5 mol of CH4 at STP
(c) 3.0 mol of H2 at STP
(d) 2.5 mol of Ne at STP
(a) 33.6L of CO2 at STP
(b) 78.4L of CH4 at STP
(c) 67.2L of H2 at STP
(d) 56.0L of Ne at STP

22. Moles of Compounds
Coefficients tell you how many moles of the compound (the whole thing) there are.
(Ex) 2 H2O  2H2 +O2
How many moles of water in the above reaction?
How many moles of H2?
How many moles of O2?
What is the mole ratio of O2 to H2? 2 2 1
1:2

23. Ratios of moles in reactions:
(ex) Given the reaction:
CH4(g) + 2O2(g)  CO2(g) + H2O(g)
How many moles of oxygen are needed for the complete combustion of 7.0 moles of CH4?
4Al(s) + 3O2(g)  2Al2O3(s)
How many moles of Al2O3 will be formed when 7.0 mol of Al react completely with O2?

24. More practice (ex) Given the reaction: N2(g) + 3H2(g)  2NH3(g)
Calculate the number of moles of NH3 produced when 2.50 moles of H2 are completely reacted.
CH4(g) + 2O2(g)  CO2(g) + H2O(g)
Calculate how many moles of CO2 will be produced when 5.93 moles of O2 is completely consumed.

25. Now all together!!  Complete the reaction, balance, and solve.
HNO3 (aq) + Mg (s) 
If g of HNO3 are reacted with Mg, how many liters of hydrogen gas will be produced?

26. Percent Composition Percent Composition:
the part of the whole compound is made up of a certain element
Found on Table T
What percent of the people in this room have brown eyes?
(Ex) Find the percent composition of potassium in potassium sulfide?
(Ex2) Find the percent composition of nitrogen in ammonium nitrate?
(Ex3) Which of the following compounds has the greatest percent composition by mass, of oxygen: BaO, CO, Al2O3, or NaClO?
Part
Whole
% =
x 100
K2S
= 2(39) / 110 x 100
= % K
NH4NO3
= 2(14) / 80 x 100
= % N
BaO:
= 16 / 153 x 100
CO:
= 16 / 28 x 100
Al2O3:
= 3(16) / 102 x 100
NaClO:
= 16 / 74 x 100
= % O
= % O
= % O
= % O

27. Percent Hydrate Hydrate:
an ionic compound that has water trapped inside the crystal lattice
Anhydrate:
a hydrate that has had the water driven off (evaporated)
The water is trapped inside the ionic crystal by molecule-ion forces of attraction. The • shows that the water is NOT bonded to the ionic compound; it is just a force of attraction.
Find the GFM of each hydrate:
(ex) CuSO4 • 5H2O
(ex2) BaCl2 • 2H2O
(ex3) BaSO4 • 7H2O
For each of the examples above find the % H2O of hydration
= (16) + 5(2 + 16)
= 250 g
= (35) + 2(18)
= 243 g
= (16) + 7(18)
= 359 g
= 5(18) / 250 x 100
= 36.00%
= 2(18) / 243 x 100
= 14.81%
= 7(18) / 359 x 100
= 35.10%

28. Finding % Water in a Hydrate in Lab
If we wanted to remove the water from the ionic compound, we could heat up the hydrate until the molecule-ion forces were broken.
In lab, how do we know all the water has been driven off (evaporated)?
When we get the same mass 3x in a row.
(ex) We have 0.5 mol of CuSO4 5H2O. How many moles of water are present?

29. % water in hydrate Lab Data:
(ex1) A hydrate has a mass of 45.3g. After being heated to a constant mass, the substance has a mass of 39.8g. What is the percent water of hydration?
(ex2) crucible g
crucible and hydrate g
crucible and anhydrate g
Given the lab data above, find the percent water of hydration.

30. Writing Empirical Formulas
The most reduced ratio of elements (moles) in a chemical formula
Ionic compounds are ALWAYS written as empirical formulas.
(ex) Na2O is an empirical formula because…
(ex2) C8H18 is not an empirical formula because…
What would be the empirical formula for C8H18?
Some molecular formulas are empirical. (ex) H2O
Sometimes it is important to know the most reduced ratio of elements in a compound (even in molecular compounds)…Do I need twice as much H as O to make a reaction work or three times as much H than O?
It is reduced (ionic)
It is not reduced (molecular)
C4H9

31. Empirical Formulas: Steps
The steps to writing an empirical formula:
Divide by the mass on the PT to get moles.
Simplify the mole ratio by dividing by the smallest #.
(ex) The percent composition ofa compound is 11.1% hydrogen and 88.9% oxygen. Write the empirical formula. =
2. =
11.1g / 1g per mol
88.9g / 16g per mol
Moles = g/GFM
= 11.1 mol H
= 5.56 mol O
11.1mol / 5.56 mol
5.56 mol / 5.56 mol
= 2 mol H
= 1 mol O
H2O

32. Empirical Formulas: Practice
(ex2) Determine the empirical formula for the compound which when analyzed showed 39.7% sodium by mass and 60.3% chlorine by mass. =
2. =
39.7g / 23g per mol
60.3g / 35g per mol
Moles = g/GFM
= 1.73 mol Na
= 1.73 mol Cl
1.73mol / 1.73 mol
1.73 mol / 1.73 mol
= 1 mol Na
= 1 mol Cl
NaCl

33. Empirical Formulas: Step 4
Step 3: If you don’t end up within a tenth of a whole number, multiply all #’s by a common factor to get whole #’s.
(ex3) A compound of zinc and phosphorus, when analyzed, showed 76.0%Zn and 24.0% P by mass. Calculate the simplest formula for the compound. =
76.0g / 65g per mol
24.0g / 31g per mol
Mol = g/GFM
= 1.17 mol Zn
= 0.77 mol P
1.17mol / 0.77 mol
0.77 mol / 0.77 mol
= 1.5 mol Zn
= 1 mol P
Multiply both numbers by 2 to get whole number ratio
= 3 mol Zn
= 2 mol P
Zn3P2

34. Molecular Formulas Molecular (true) Formula:
gives the actual # of moles of elements in a molecule (not reduced)
**Only molecular compounds have molecular formulas.
(ex) C6H12O6

35. Molecular Formulas: Steps
The steps to writing a molecular formula:
Find the empirical formula
Find the mass of the empirical formula
Find how many times larger the GFM is than the empirical formula and multiply the subscripts by that number
(ex) What is the molecular formula of a compound that has a GFM of 92g and an empirical formula of NO2? =
NO2
14 + 2(16)
= 46g
92g / 46g
= 2 times as larges as the empirical formula
N2O4

36. Molecular Formulas: Practice
(ex2) A compound of phosphorus and oxygen, when analyzed, showed 39.24% P and 60.76% O by mass. Calculate the simplest formula for the compound. =
Find the molecular formula if the GFM of this compound is 158g.
= mass of empirical
formula
39.24g P g O
39.24g / 31g per mol
60.76g / 16g per mol
Mol = g/GFM
= 1.27 mol P
= 3.80 mol O
1.27mol / 1.27 mol
3.80 mol / 1.27 mol
= 1 mol P
= 3 mol O
PO3
= (16)
= 79g
= 158g / 79g
= 2
P2O6

37. (ex) Bill Nye measures out 62g of NaNO3
(ex) Bill Nye measures out 62g of NaNO3. How many moles of NaNO3 does Bill have? (ex) You are asked to measure out 1.49 moles of SCl2. How many grams will you need to measure out on the scale?