1. Unit 11: Chemical Reactions
Chapter 10 in text

2. Chemical Formulas 1. Empirical formula 2. Molecular formula
Identifies elements present in terms of simplest whole number _________
Examples: table salt, NaCl; water, ________
2. Molecular formula
Identifies ________ number of atoms in a molecular compound
Example: water, H2O; Glucose ____________
Not table salt, NaCl (ionic compound)
ratio
H2O
exact
C6H12O6

3. Chemical Formulas 3. Structural formula
Represents ________________ of atoms within a molecule
Related to _______ structure of molecule
arrangement
3 – D

5. Empirical or molecular formula?
Ionic - lacking discrete unit, or molecule
Composed of __________ and nonmetallic elements
Electronegativity difference > _______
Molecular
____________ compounds
Usually nonmetals bound to _______________
Molecular and empirical formulas can be different
Glucose – molecular C6H12O6 vs. empirical _____________
metallic
1.7
Covalent
other nonmetals
CH2O

6. Molecular & Formula Weights
___________ of atomic weights of all atoms in chemical formula
Molecular weight
Formula weight of a _______________ substance
Term often used for non-molecular (________) substances
Total sum
molecular
ionic

8. Percent composition of compounds
Finding the mass percentage of an individual _________ from the formula weight
element

11. Calculate Empirical Formula
Divide the percent composition or mass of each element by their ______________.
Divide the number of each element you just found by the ___________ number you just found.
If you have a ________ such as 0.5, multiply all quotients by 2
0.25 multiply by 4, and 0.33 multiply by 3.
The answers will be the ____________ of the empirical formula.
The empirical formula should be the simplest __________ number ratio of the subscripts.
atomic mass
smallest
decimal
subscripts
whole

12. Empirical Formulas Divide each mass by the atomic mass
2.3g Na / 23 = __________
0.8 g O/ 16 = ___________
Divide the quotient of each element by the smallest quotient.
Na /_______ = _____
O /______ = _____
The products will be the subscripts of the empirical formula.
___________
0.1
0.05 2
0.05
0.05 1
Na2O

13. Empirical Formulas Divide each mass percentage by the mass number
Divide the quotient of each element by the smallest quotient.
S /______ = _______
O / _______ = ______
The products will be the subscripts of the empirical formula.
__________
1.56
3.125
1.56 1
1.56 2
SO2 Sulfur dioxide

14. Chemical reactions
Occur through formation and ___________ of chemical bonds between atoms
Involve changes in matter, creation of __________ materials and __________ exchanges
Chemical equations - concise _______________ of chemical reactions
breaking
new
energy
representations

15. Chemical equations Reactants - substances existing _________ reaction
Products - substances existing __________ reaction
Word representation only is not precise
Chemical symbols and formulas needed for ____________ purposes
before
after
quantitative

17. Balancing equations
Law of conservation of mass - atoms are neither ________ nor destroyed in chemical reactions
Mass of the products = Mass of reactants
No mass defect, like with nuclear reactions
Any _________ released was stored in the bonds!
Change ____________ in front of chemical formulas, not ___________ within formulas, to balance
created
energy
coefficient
subscripts

20. Types of Reactions Decomposition reactions Combination reactions
Replacement reactions
- Single ion replaces another ion in a compound
Ion exchange reactions
- Two cations switch places in their compounds

22. AB + CD  AD + CB
Positive ions change/swap which anion they are bonded with
(i.e. A was with B and A is now with D)

23. Types of chemical reactions
Oxidation-reduction (redox) reactions
Oxidation - ___________ of electrons by an atom
Reduction - ___________ of electrons by an atom
__________ agents - substances taking electrons from other substances (oxygen, chlorine)
_______________ agents - supply electrons to oxidizing agents (hydrogen, carbon)
loss
gain
Oxidizing
Reducing

25. Chemical Reactions
Atoms are conserved – Atoms do not get used up in a __________ reaction
Burning a piece of paper is a combustion reaction
The atoms in the paper and the __________ in the air react to form the ash, smoke, and the CO2 that exist after the reaction occurs
Mass is ____________ – Same number and type of atoms, same mass is present
Law of combining volumes (gases)
Gases at the same temperature and pressure contain equal numbers of molecules
chemical
oxygen
conserved

28. Units of Measurement Used with Chemical Equations
Atomic mass unit = 1 a.m.u. (1 u)
Roughly the mass of one _________ or one ___________
1 u = x kg
Defined as 1/12 mass of carbon-12 ______
The Periodic Table lists the atomic masses of each element in ______
Since we deal with more than one atom at a time, we need a way to go from the macroscopic scale to the atomic scale and vice-versa…
proton
neutron
atom
a.m.u.

29. The Mole
The Mole – one mole of carbon-12 has a mass of 12 grams and contains 6.02 x 1023 atoms of carbon
A mole is a number of something, like a dozen means ____ of something or a gross is _____ of something – Could you have a mole of cars? ______
6.02 x 1023 is known as Avogadro’s number and is the number of things in a mole of that thing
One mole of cars would be ____________ cars
A mole of carbon-12 atoms is defined as having a mass of _____________ 12
144
Yes!
6.02 x 1023
12 grams

30. Molar weights
____________ atomic weight - mass in ___________ equal to atomic weight
______________ formula weight - mass in grams equal to formula weight (ionic)
Gram molecular weight - mass in grams equal to molecular weight of the _____________
Gram
grams
Gram
compound

32. Quantitative Use of Equations
Possible interpretations:
Molecular ratios of reactants and products
Mole ratios of reactants and products
Mass ratios of reactants and products

34. Stoichiometry
The study of the amount of reactants needed and products generated in a _________ _________.
There are _________ steps needed to solve a Stoichiometry problem.
If your initial units are __________ and your final units are moles, skip steps ___ and ___.
You only need the mole-mole ratio from the coefficients of the chemical reaction.
chemical
reaction
five
moles 2 4

35. Stoichiometry Write a complete and _____________ chemical equation.
Convert quantity of given into __________.
Multiply by the mole _____ taken from the ________________ in the chemical equation.
Convert from moles into the desired unit.
_____________ units to ensure that correct conversion factors were used.
balanced
moles
ratio
coefficients
Cancel out

36. Stoichiometry
What volume of ammonia is produced if 25.0g of nitrogen (diatomic molecule) is reacted with excess hydrogen?
Write a complete and _____________ chemical equation.
_______________________________________
Convert mass of nitrogen into moles by dividing by the molar mass.
balanced
N2 + 3H2  2NH3
x 1mol N2
28 g N2
25.0 g N2

37. Stoichiometry Multiply by the mole ratio.
________________________________________
Convert moles into the desired units
Cancel units.
Answer: _______________
x 1mol N2
28 g N2
x 2 mol NH3
1 mol N2
25.0 g N2
x 1mol N2
28 g N2
x 2 mol NH3
1 mol N2
x 17.0 g NH3
1 mol NH3
25.0 g N2
x 1mol N2
28 g N2
x 2 mol NH3
1 mol N2
x 17.0 g NH3
1 mol NH3
25.0 g N2
30.4 g NH3

38. Limiting Reagents
A limiting reagent is a reactant that is in less molar quantity or the reactant that determines the quantity of _________________.
For example, if we react 2.5 grams of sodium with 300,000 grams of chlorine, we wouldn’t get 300,002.5 grams of sodium chloride
We would run out of sodium atoms before we ran out of chlorine atoms
Therefore, the sodium “limits” the amount of sodium chloride that can be produced and is our limiting reagent
product produced

39. Limiting Reagents
Problem: How many grams of NH3 can be produced from the reaction of 20 g of N2 and 6 g of H2?
Calculate the moles of each reactant.
_______________________________________
20 g N2
x 1 mol N2
28 g N2
= mol N2
x 1 mol H2
2 g H2
6 g H2
= 3 mol H2

40. Limiting Reagents According to the equation, N2 + 3H2  2NH3
For every 1 mole of nitrogen, I would need ________ moles of hydrogen.
However, I have been given moles of N2 and 3 moles of H2. To use all the N2, I need 3 x = 2.14 moles of H2.
There would be __________________ moles H2 remaining.
Therefore, nitrogen is the ___________ reagent and hydrogen is the __________ reagent.
three
3 – 2.14 = 0.86
limiting
excess

41. Limiting Reagents
This means my Stoichiometry must be based on the quantity of nitrogen since I will need all of it producing the Ammonium.
_______________________________________
0.86 mol H2 x 2 g H2/1 mol H2 = 1.72 g H2 remains
How much hydrogen was used?
Answer: ______________________
x 1mol N2
28 g N2
x 2 mol NH3
1 mol N2
X 17.0 g NH3
1 mol NH3
20.0 g N2
= g NH3
6 g H2 – 1.72 g H2 = 4.28 g H2