1. Chemical Quantities

2. Stoichiometry
The relationship between the relative quantities of substances taking part in a reaction or forming a compound

3. Measuring Matter Matter can be measured by 3 methods By count
the number of CDs you have
The number of atoms or molecules you have
By mass or weight
You buy vegetables by the pound
The mass(g) of a compound or element you have
By volume
You buy soda in two liter bottles
The number of milliliters of a substance you have

4. Finding Mass from a Count
What is the mass of 90 average-sized apples if a dozen apples has a mass of 2.0kg?
1 dozen
2.0kg
90 apples
12 apples
1 dozen apples
= 15 kg or 15,000g

5. When determining quantities in chemistry the method is the same…..
You just use different units !!!!!!!!

6. Chemical Units

7. The Mole
Mole = 6.02 x particles of a substance – could be atoms, molecules, or formula units
Ex: A mole of hydrogen atoms is 6.02 x hydrogen atoms, a mole of paper is 6.02 x piece of paper
6.02 x is known as Avogadro’s number- Avogadro was an Italian scientist in the 1800’s helped clarify the difference between atoms and molecules

8. A mole of any substance contains 6.02 x 1023 representative particles

9. mole is a unit
TedEd

10. How many atoms of carbon would there be in one mole?
How many atoms of silver would there be in one mole?
How many atoms of gold would there be in a mole?
How many molecules of water would there be in a mole of water?

11. The Mass of a Mole (MOM) of an element = molar mass in grams

12. Mass of a Mole
The atomic mass of an element expressed in grams is the mass of one mole of that element.
Example- 1 mole of carbon is 12.0 grams.
The molar mass of any two elements contain the same number of atoms

13. Example Ex: a dozen apples – 12 apples a dozen oranges – 12 oranges
a mole of carbon – 12.0 g – 6.02 x 1023 atoms
a mole of lithium – 6.94 g – 6.02 x 1023 atoms

14. Mass of a Mole of a Compound
Find the number of grams of each element in the compound, then add the masses together.
Ex: Find the molar mass of water.
Step 1- write down the formula H2O
Step 2- determine masses of each element
2 H atoms x 1.0 g = 2.0 g H
1 O atom x g = 16.0 g O
Step 3- add masses together 16.0 g O g H = 18.0 g H2O
The molar mass of water is 18.0 g

15. Converting Moles to Particles
formula

16. Converting Particles to Moles
formula

17. Mole to Volume STP- Standard Temperature and Pressure
0 degrees Celsius and 1 atmosphere (101.3 kPa)
At STP 1 mol (6.02 x 10 23) any gas occupies a volume of Liters
Calculating Volume at STP
Calculating molar mass from density

18. Ex: sulfur dioxide is a gas produced by burning coal
Ex: sulfur dioxide is a gas produced by burning coal. It is an air pollutant and one of the causes of acid rain. Determine the volume, in liters, of moles of sulfur dioxide gas at STP.

19. What does an equation represent?
2 Na + Cl2  2 NaCl or
2 Na + Cl2  2 NaCl

20. Mole Ratios
2H2 + O H20
if you produced 3.65moles of water, how many moles of oxygen would have been used?

21. Mole Ratios
Mole ratios can be used to calculate the moles of one chemical from the given amount of a different chemical
Example: How many moles of chlorine is needed to react with 5 moles of sodium (without any sodium left over)?
2 Na + Cl2  2 NaCl
5 moles Na 1 mol Cl2
2 mol Na
= 2.5 moles Cl2

22. Stoichiometry STEP 1: You MUST start with a balanced equation
STEP 2:Start with known amount.
STEP 3: Use the mole ratio as a conversion factor

23. Practice
How many moles of oxygen are produced by the decomposition of 4.7 moles of potassium chlorate?

24. Mole ↔Mass Conversions
Most of the time in chemistry, the amounts are given in grams instead of moles
We still go through moles and use the mole ratio, but now we also use molar mass to get to grams

25. MOLES moles use balanced equation)
MOLES  GRAMS GRAMS  MOLES
CALCULATE MOLAR MASS

26. Practice
Calculate how many moles of oxygen are required to make 10.0 grams of aluminum oxide

27. Now add one more step…
Ex. Calculate how many grams of ammonia are produced when you react 2.00 grams of nitrogen with excess hydrogen.
N2 + 3 H2  2 NH3

28. Practice
How many grams of calcium nitride are produced when 2.00 grams of calcium reacts with an excess of nitrogen?

29. Solution Chemistry

30. solutions Homogeneous mixture = solution
A solution contains a solute (the thing being dissolved) and a solvent (the thing doing the dissolving)
Usually refer to as “dilute” or “concentrated”
More specifically-The concentration of a solution is a ratio of solute to solvent (usually)

31. Concentration can be expressed in a variety of ways
Mass/mass percent (m/m)
Mass/volume percent (m/v)
Volume/volume percent (v/v)
Molarity (moles/L)
Molality (moles/kg)

32. Remember percent…. The percent of anything = part of interest x 100%
whole thing
for solutions – part of interest is usually the solute and the whole is usually the solvent

33. mass/mass percent concentration
Mass of solute x 100 % = m/m % mass of solution Mass of solute + mass of solvent

34. mass/mass percent concentration
If a student prepares a solution from 5.00 g of NaCl dissolved in g of water, what is the mass/mass percent concentration?

35. As a unit factor… 100 g of NaCl (95 g water and 5 g NaCl)
If the sodium solution is a 5% solution, that means that for every 5 grams of solute there is 100 g of solution
5 g NaCl x 100 = 5%
100 g of NaCl
(95 g water and 5 g NaCl)

36. Ex:
A glucose solution is 5%, what is the mass of a solution that contains 1.25 g of glucose?
( 5 g glucose / 100 g solution)
1.25 g glucose x g solution = 25 g glucose
5 g glucose

37. Molar concentration (moles/Liter)
molarity (M) = moles solute
Liter solution

38. Ex:
A drain cleaner contains g of NaOH dissolved in 1.00 L of solution, what is the molarity of the solution?

39. Dilution
The amount of solute does not change during dilution M1V1 = M2V2

40. Dilution
What volume of a 6.00 M solution do I need to make 5.00 L of 0.1 M solution ?
6 M x V1 = 0.1 M x L