Chapter 2
Isotopes
 Isotopes occur when different numbers of neutrons
are in the nucleus, but all else remains the same (# of
protons, electrons, and atomic number).
 Recall – Isotope notation
 Ex:
Average Atomic Mass
 Isotopes are always present in the same amounts.
 Ex. The isotopes for Magnesium are 79% Mg-24, 10%
Mg-25, and 11% Mg-26.
 The average atomic mass (found on periodic table) is
a calculated value taking into account the masses of all
the elements isotopes.
 Do Sample Problem on Page 45
 Do Practice Problems on Page 45 #s 1-3.
 Do Section Review on Page 46 #s 1-5 (a only)
Section 2.2
 In chemistry we need to know how many atoms,
formula units or molecules are present in a sample.
How many oxygen atoms are in a 10g, 20 g, or 30 g
sample?
 The actual number is too large to count. So we group
numbers to make it easier to handle.
Ex: dozen, pair, score, gross, ream
Mole
 The mole is defined as the amount of substance that
contains as many particles, atoms, molecules, or
formula units as exactly 12 g of carbon-12.
 Particles is a general term covering atoms, molecules,
or formula units.
 Atoms – you may have to identify the number of atom
overall in a compound or the number of one particular
atom.
 Molecules result from covalent bonding.
 Formula units result from ionic bonding.
 One mole of a substance contains
6.02214199 × 1023 particles.
 Avogadro’s Constant (NA) is
6.02214199 × 1023 ____/mol.
Conversions
 Use conversion factors to solve the following problems.
 How many eggs in 10 dozen?
 One box of paper contains 8 reams, how many sheets of
paper?
 How many aluminum atoms in 2 moles of aluminum?
Converting Moles to Number of
Particles
 We can convert moles to/from number of atoms,
molecules, and formula units.
 1 mole of potassium (K) is 6.02 × 1023 atoms
 1 mole of potassium sulfide (K2S) is 6.02 × 1023 formula
units
 1 mole of sugar (C6H12O6) is 6.02 × 1023 molecules
 The formula N = n × NA may help.
 Where N = number of particles.
n = amount (moles)
NA = Avogadro’s Constant (6.02 × 1023)
 Sample Problem on Page 51
 Practice Problems on Page 51-52 #s 5-11
Converting Number of Particles to
Moles
 You can convert number of particles to moles using
conversion factors or by rearranging the formula.
 N = n × NA
 n = N/NA
 See Sample Problem on Page 53
 Do Practice Problems on Page 53 #s 12-15
 Do Section Review on Page 54 #s 1-8 and 10
Molar Mass
Section 2.3
 1 mole of Carbon-12 atoms (6.02 × 1023 atoms) has a
mass of 12 g.
 1 mole of any element has a mass in grams that is
equivalent to the element’s average atomic mass.
 The mass of 1 mole of any substance is its molar mass
(M) in g/mol.
 To find the molar mass of an element, read it off the
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periodic table.
To find the molar mass of a compound, find the molar
mass of each element, multiply by the number of
atoms, and then add up all the mass values.
Ex. 1. Find the molar mass of SO3.
Ex. 2. Find the molar mass of potassium sulfide.
Do Sample Problem on Page 56
Do Practice Problems on Page 57 #s 16-19
Convert from Number of Particles
to Mass
 6.02 × 1023 atoms of carbon = 1 mol of carbon = 12.0 g
carbon
 6.02 × 1023 molecules of CO2 = 1 mol CO2 = 44.0 g CO2
 Convert 1.204 × 1023 atoms of Mg to grams.
 See Figure 2.9 on Page 57
Converting from Moles to Mass
 How many grams are in 10.0 mol of dinitrogen
trioxide?
Get the molecular formula – N2O3
2. Get the molar mass of each atom and the molar mass
of the molecule.
3. Use conversion factors (pick appropriate CF)
4. Solve
1.
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Sample Problem on Page 58
Do Practice Problems on Page 59 #s 20-23
Converting from Mass to Moles
 How many moles of N2O3 are in 1520.4 g?
 Sample Problem on Page 60
 Do Practice Problems on Page 60 #s 24-27
 Worksheet
Converting Between Moles, Mass,
and Number of Particles
 To do multistep problems be sure to identify what you
are beginning with, what you need, and the path to get
there.
 See Sample Problem (Particles to Mass) on Page 62
 See Sample Problem (Mass to Particles) on Page 63
 Do Practice Problems on Page 63 #s 28-33 and on Page
64 #s 34-37.
Molar Volume
Section 2.4 Volume of a Gas
 Air bags, balloons, basketballs, and tires all contain
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gas.
What happens if you put too much air in a balloon?
It explodes. Why? Too much pressure.
The gases are always moving inside the object,
colliding with each other and the inside of the balloon.
Pressure is the force exerted on an object per unit of
surface area (Pa or kPa).
For an air bag to inflate, would you want a lot of
pressure or just a little?
 The volume of a gas is indefinite – meaning you can
increase or decrease the volume by adjusting the size
of the container.
 Decreasing the size, puts the particles in a smaller space,
increasing the collisions, speed, and pressure.
 Increasing the size, decreases the pressure because there
is more space and fewer collisions.
 The volume of a gas is also affected by temperature.
 Increasing the temperature increases the volume of
the gas.
 The gas particles speed up, move faster, and have more
collisions.
 Decreasing the temperature decreases the volume of
the gas.
 Ex. Put a balloon in the freezer and it will shrink.
 When the temperature falls the gas particles slow down
and there are fewer collisions.
 Pressure and volume are inversely proportional.
 Temperature and volume are directly proportional.
 Standard conditions for gases:
 Standard pressure (101.3 kPa) is the average atmospheric
pressure at sea level.
 Standard temperature is 0oC, (the freezing point of
water) or 273 K.
 Therefore STP is 0oC and 101.3 kPa
Molar Volume of Gases – The
Beginning
 French scientist, Joseph Louis Gay-Lussac, devised the
law of combining volumes: when gases react, the
volume of the reactants and products, measured at
equal temperatures and pressures, are always whole
number ratios.
 Around the same time, John Dalton, found that the
masses of elements that combine to form compounds
can be expressed as small whole numbers.
 Avogadro realized that he could relate the volume of a
gas to the amount that is present (calculated from the
mass).
 The number of gas molecules corresponded to the
volume ratios.
 Avogadro’s hypothesis – equal volumes of all ideal
gases at the same STP contain the same number of
molecules.
 This was the beginning of the mole concepts that
connects the amount of a substance to its mass.
 Thought Lab Page 69
Standard Molar Volume
 The molar volume of a gas at STP is 22.4L/mol.
 This fact is true for any ideal gas.
 The particles of an ideal gas have zero volume and do
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not attract each other. Is this possible?
Define matter. Matter is anything that has mass and
takes up space.
The ideal gas is a myth. It does not exist.
Particles do have volume and can attract each other.
The molar volume of real gases vary slightly around
the accepted value of 22.4 L/mol.
 See Table 2.3 on Page 70.
 For all your calculations, 1 mol = 22.4 L is all you need.
 Example:
Find the volume, in litres, of 3.76 mol of chlorine
gas at STP.
 See Sample Problems on Pages 70 and 71.
 Do Practice Problems on Page 73 #s 38-43.
 Chapter 2 Review Page 75-76 #s 2, 7, 9, 10-15, 17-19, 21-24,
26, & 27