1. The Mole: Avogadro’s number

2. How much is: A dozen? 12 A century? 100 A mole? 6.02 x 1023
(602,000,000,000,000,000,000,000)

3. Can you count a mole of pennies?
If you could count 5 per second, it would take you
6.02 x ÷ 5 pennies/second ÷ 60 sec/min ÷ 60 min/hr ÷ 24hrs/day ÷ 365 days/yr =
3,800,000,000,000,000 years!

4. Can you spend a mole of dollars?
If you could spend $1,000,000 every second it would take you
6.02 x ÷ $1,000,000/sec ÷ 60 sec/min ÷ 60 min/hr ÷ 24hrs/day ÷ 365 days/yr =
19,000,000,000 years!

5. The Mole Map    x x Mole x Mass # Particles 22.4L Gas Volume @ STP
Molar
Mass x
Mass

6. When measuring amounts, you can count or you can mass them.
If I want 2 dozen baseballs, I can count 24 baseballs
Or I can mass 16 kg of baseballs.

7. How many tennis balls are in 6 kg?24
( 2 dozen)

8. How many tennis balls are in a mole?
6.02 x 10^23
in a mole
12 in a dozen

9. Since we can’t count a mole of atoms, we MUST mass chemicals to measure moles
6.02 x atoms of sulfur
32.07 grams of sulfur
6.02 x atoms of carbon
12.01 grams of carbon

10. How do we measure moles?
mole = number of particles equal to the number of atoms in 12 g of C-12
1 atom of C-12 weighs exactly 12 amu
1 mole of C-12 weighs exactly 12 g
The number of particles in 1 mole is called Avogadro’s Number = x 1023
1 mole of C atoms weighs g and has x 1023 atoms
the average mass of a C atom is amu

11. How do we measure moles?
The atomic mass on your periodic table is the mass of a mole of atoms of that element.
What is the mass of a mole of copper atoms?
63.55 g
So, to count 6.02 x 1023 copper atoms, we mass out g on the scale.

12. Mole and Mass Relationships
Sulfur
32.06 g
1 mole
Carbon
12.01 g

13. Find the mass of: A mole of silicon atoms 28.09 g
6.02 x 1023 atoms of nitrogen
14.01 g
6.02 x 1023 atoms of sodium
22.99g
2 moles of sodium atoms
45.98 g

14. How many atoms are in: A mole of silicon 6.02 x 1023
14.01 g of nitrogen
2 moles of sodium
12.04 x 1023
45.98 g of sodium

15. How many things are in: A mole of footballs 6.02 x 1023 footballs
A mole of water
6.02 x 1023 molecules
2 moles of pencils
12.04 x 1023
½ mole of lead
3.01 x 1023 atoms

16. Molar mass is the mass of one mole of a substance
To find the molar mass of an element, look on the periodic table.
To find the molar mass of a compound, add all the masses of its elements

17. Chemical Formulas as Conversion Factors
1 spider  8 legs
1 chair  4 legs
1 H2O molecule  2 H atoms  1 O atom

18. Molar Mass of Compounds
the relative weights of molecules can be calculated from atomic weights
Formula Mass = 1 molecule of H2O
= 2(1.0 amu H) amu O = 18.0 amu
since 1 mole of H2O contains 2 moles of H and 1 mole of O
Molar Mass = 1 mole H2O
= 2(1.01 g H) g O = g

19. Find the molar mass of: Ammonium phosphate NH4+ PO43- (NH4)3PO4=
=149.12g/mol
Carbon dioxide
CO2 =
= g/mol

20. Find the molar mass of: Hydrogen gas Elemental hydrogen H2
= 2.02 g/mol
Elemental hydrogen H
= 1.01 g/mol

21. Find the molar mass of: Iron Iron (III) hydroxide Fe = 55.85 g/mol
Fe3+ OH-
Fe(OH)3 =
= g/mol

22. Converting to and from moles.
Converting between moles and mass requires the molar mass of the substance from the periodic table.
Element: Ag = g/mol
Ionic compound: CaCl2 = g/mol
Covalent compound: NO2 = g/mol
Always keep at least two decimal places on all values taken from the periodic table.

23. Converting to and from moles.
To convert from moles to grams, multiply by molar mass:
0.500 mol H2O x (18.02g/mol) = 9.01g H2O
To convert from grams to moles, divide by molar mass:
54g H2O x (1mol/18.02g) = 3.0 mol H2O

24. Converting to and from moles.
For gases, use the fact that at STP, 1 mol of any gas has a volume of 22.4 Liters.
What is STP? Standard Temperature and Pressure
Standard Temperature = 273K or 0.0°C
Standard Pressure = 1 atmosphere = 760 mm Hg (barometric) = kPa.

25. Converting to and from moles.
To go from moles to volume, multiply by 22.4L.
3.00 mol x (22.4L/mol) = 67.2L of gas
To go from volume to moles, divide by 22.4L
44.8L x (1mol/22.4L) = 2.00 moles of gas

26. Converting to and from moles.
To convert between moles and particles, simply multiply or divide by Avogadro’s number.
2.0 mol x (6.02 x 1023 particles/mol) = 1.2 x 1024 particles
3.1 x 1024 particles x (1 mol/ 6.02 x 1023 particles) = 5.0 mol

27. Remember unit factors?

28. Converting to and from moles.
A convenient tool for making these conversions is called a “mole map.”
With the mole at the center, we can put all of the aforementioned calculations together into one simple picture.

29. The Mole Map    x x Mole x Mass # Particles 22.4L Gas Volume @ STP
Molar
Mass x
Mass

30. Percent Composition Percentage of each element in a compound By mass
Can be determined from
the formula of the compound
the experimental mass analysis of the compound
the total mass of each element
The percentages may not always total to 100% due to rounding

31. What percentage of water is Oxygen?
Formula of the compound
* H2O
Mass of the compound
* 2 (1.01 g/mol) g/mol = g/mol
Mass of each element
* H is 1.01 g/mol, O is g/mol

32. Mass Percent as a Conversion Factor
the mass percent tells you the mass of a constituent element in g of the compound
the fact that NaCl is 39.0% Na by mass means that 100.0g of NaCl contains 39.0g Na
this can be used as a conversion factor
100.0 g NaCl  39.0 g Na

33. Empirical Formulas
The simplest, whole-number ratio of atoms in a molecule is called the Empirical Formula
can be determined from percent composition or combining masses
The Molecular Formula is a multiple of the Empirical Formula
% A mass A (g) moles A
100g MMA
% B mass B (g) moles B
100g MMB
moles A
moles B

34. Empirical Formulas Hydrogen Peroxide Molecular Formula = H2O2
Empirical Formula = HO

35. Benzene
Molecular Formula = C6H6
Empirical Formula = CH
Glucose
Molecular Formula = C6H12O6
Empirical Formula = CH2O

36. Finding an Empirical Formula
convert the percentages to grams
skip if already grams
convert grams to moles
use molar mass of each element
divide all by smallest number of moles
round or multiply all mole ratios by number to make all whole numbers
if ratio ?.5, multiply all by 2; if ratio ?.33 or ?.67, multiply all by 3, etc.
skip if already whole numbers

37. Determine the empirical formula of a compound containing 80
Determine the empirical formula of a compound containing 80.0 grams of carbon and 20.0 grams hydrogen.

38. Grams to moles

39. Divide by smallest

40. Write Empirical Formula
CH3

41. All these molecules have the same Empirical Formula
All these molecules have the same Empirical Formula. How are the molecules different?
Name
Molecular
Formula
Empirical
glyceraldehyde
C3H6O3
CH2O
erythrose
C4H8O4
arabinose
C5H10O5
glucose
C6H12O6

42. All these molecules have the same Empirical Formula
All these molecules have the same Empirical Formula. How are the molecules different?
Name
Molecular
Formula
Empirical
Molar
Mass, g
glyceraldehyde
C3H6O3
CH2O
90.09
erythrose
C4H8O4
120.12
arabinose
C5H10O5
150.15
glucose
C6H12O6
180.18

43. Molecular Formulas
The molecular formula is a multiple of the empirical formula
To determine the molecular formula you need to know the empirical formula and the molar mass of the compound
Molar Massreal formula
Molar Massempirical formula
= factor used to multiply subscripts

44. What is the molecular formula for ethane if it has a molar mass of 30
What is the molecular formula for ethane if it has a molar mass of g/mol?
CH3= g/mol
Molecular formula = 2 x the empirical formula
C2H6

45. Determine the Molecular Formula of Cadinene if it has a molar mass of 204 g and an empirical formula of C5H8

46. Solutes and Solvents Solution: a homogenous mixture
Solution: a homogenous mixture
Solute: thing that dissolves
Solvent: thing that does the dissolving (found in the largest amounts)
If the solvent is water, then it is called an aqueous solution

47. Solubility Why does sugar “disappear” in your iced tea?
Why does sugar “disappear” in your iced tea?
How do fish breathe underwater?
Why does soda go flat faster when left out than when it is refrigerated?
It is all based on solubility!

48. Solubility
Example: iced tea
Solute
sugar
tea
Solvent
water

49. States and Solutions Solutions can be any state of matter
Solutions can be any state of matter
Solid-solid: alloys (gold jewelry, brass, etc.)
Solid-liquid: salt water, sugar water
Liquid-liquid: vinegar, peroxide, rubbing alcohol
Liquid-gas: soda, champagne, O2 in H2O
Gas-gas: air, air tanks (scuba)

50. How Things Dissolve
Need to find/ create “holes” in water for the dissolving substance to move into
Need to over come hydrogen bonding between water (or solvent) molecules
Get interactions between water molecules and molecules of the solute
Ion-dipole interactions
Dipole-dipole (and H bonding)

51. Why some coffees “Put hair on your chest.”
“Strong” coffee has more coffee dissolved in a given amount (say 1 pot) than “weak” coffee.
Strong coffee = concentrated
Weak coffee = dilute
Concentration: the amount of solute in a given amount of solvent (or solution).

52. Molarity (M) M = moles of solute
Most common way to express concentration
Molarity is the number of moles of solute dissolved in each liter of solution
Formula
M = moles of solute
liters of solution
Dependent on temperature
The higher the molarity the stronger the concentration

53. Practice Problems
1. What is the molarity when 6.0 moles of glucose is dissolved in water to make 3.0 L of solution.
2. How many moles of sodium chloride are there in 500 mL of 4.0 M solution?
3. What is the volume of 3.0 M solution that contains 15 moles of glucose?

54. How does something so strong become so weak?
The answer is dilution.
The more dilute something is, the lower the concentration (it’s weaker).
To accomplish this, add more solvent
How do we know how much to add?
M 1V1 = M 2V2
Typically start with a highly concentrated solution and dilute down to what you need

55. Figure 15. 8: Process of making 500 mL of a 1
Figure 15.8: Process of making 500 mL of a 1.00 M acetic acid solution.

56. Mole Day is October 23!