Chapter 3.3 and 7.3 The Mole
I.
II.
III.
Measuring Matter
a. Counting Particles, weighing Mass, and measuring the Volume
b. Use conversion factors when doing dimensional analysis to change units
The Mole- the amount of a substance
a. Avogadro’s Number (counting particles)
i. 1 mole = 6.02x1023 particles (general)
ii. Particle Types
1. Atoms (specific) 1mol Al = 6.02x1023 atoms of Al
2. Molecules (specific) 1mol CO2 = 6.02x1023 molecules of CO2
3. Formula Units (specific)
1mol MgCl2 = 6.02x1023 formula units of MgCl2
b. A mole is still a very small visual (physical) quantity, even though the count
(number) is very large: 602,000,000,000,000,000,000,000 particles
c. Conversions between particles and moles
Given particles of A x
1
1mol A
6.02x1023 particles of A
= x mol A
Given moles of B
1
6.02x1023 particles of B
1mol B
= x particles B
x
Example 1: How many atoms are in 1 mole of CO2?
You now know that in 1 mole of CO2 there are 6.02x1023 molecules of
CO2 and that 1 molecule of CO2 has 3 atoms. Do the conversion…..
Example 2: How many formula units are in one mole of MgCl2?
Example 3: If all moles contain the same number of particles (6.02x1023),
would you predict that moles of two substances would have different or the
same mass? Why or why not?
The Mole- Formula Mass (grams): 1mol = 6.02x1023 particles = the mass of
that quantity of particles, which can be calculated using the periodic table.
a. Gram Atomic Mass – mass of one mole of an atom
1mol H = 6.02x1023 atoms of H = 1g of H
1mol O = 6.02x1023 atoms of O = 16g of O
1mol Na = 6.02x1023 atoms of Na = 23g of Na
1mol Cl = 6.02x1023 atoms of Cl = 35g of Cl
b. Gram Molecular Mass – mass of one mole of a molecule (NM+NM)
1mol H2O = 6.02x1023 molecules of H2O = 18g of H2O because
Hydrogen = 1g x 2atoms = 2g H
Oxygen = 16g x 1atom = 16g of O
2g of H + 16g of O = 18g of H2O
1mol CO2 = 6.02x1023 molecules of CO2= 44g of CO2 because
Carbon = 12g x 1atom = 12g of C
Oxygen = 16g x 2atoms = 32g of O
12g of C + 32g of O = 44g of CO2
IV.
V.
VI.
1mol NO2 = 6.02x1023 molecules of NO2 = 46g of NO2 because
Nitrogen = 14g x 1atom = 14g of N
Oxygen = 16g x 2atoms = 32g of O
14g of N + 32g of O = 46g of NO2
c. Gram Formula Mass – mass of 1mol of ionic comp, formula unit, (M+NM)
1mol NaCl = 6.02x1023 formula units NaCl = 58g Nacl
Sodium = 23g x 1atom = 23g Na
Chlorine = 35g x 1atom = 35g Cl
1mol CaI2 = 6.02x1023 formula units = 294g CaI2
Calcium = 40g x 1atom = 40g
Iodine = 127g x 2atoms = 254g
1mol (NH4)2CO3 = 6.02x1023 f.u. = 390g (NH4)2CO3
Nitrogen = 14g x 2atom = 28g N
Hydrogen = 1g x 8atoms = 8g H
Carbon = 12g x 1atom = 12g C
Oxygen = 16g x 3atoms = 48g C
d. Molar Mass – the mass of one mole of any particle
i. Specifically –
Atomic Mass is the mass of an atom.
Molecular mass is the mass of a molecule.
Formula mass is the mass of a formula unit.
ii. Generally – We use the term formula mass for all chemical formulas
regardless of the type of particle (atom, molecule, ionic). We also use
the term molar mass for the mass of one mole of any type of particle.
Conversions between grams and moles
Given grams of A
x
1mol A
= xmol A
1
FM of A
Given moles of B
x
FM of B
= xg B
1
1mol B
Example 1: How many grams are there in 7.35mol of glucose, C6H12O6? Start by
calculating the formula mass of glucose.
Example 2: How many moles are there in 500.g of lead (II) nitrate, Pb(NO3)2?
The Mole also equals a volume of 22.4L of a gas at STP
a. Gas – spreads out as far as possible. No attraction between the molecules.
b. STP = Standard Temperature and Pressure
i. Standard Temperature = 273K = 0ºC
ii. Standard Pressure = 101.3kPa = 1atm = 760mmHg
c. 1 Liter = 1000ml
Conversions between volume and moles
Given liters of A
x
1mol A
= x mol A
1
22.4L of A
Given mol of B
1
x
liters of B
1mol B
=xLB