By definition, the oxidation number of an atom is the charge that atom would have if the compound was composed of ions. It is a way of keeping track of where the electrons go in reactions.

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Chapter 10

Chemical Quantities

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How you measure how much?

• You can measure mass,
• or volume,
• or you can count pieces.
• We measure mass in grams.
• We measure volume in liters.
• We count pieces in MOLES.

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Moles

• Defined as the number of carbon atoms in exactly 12 grams of carbon-12.
• 1 mole is 6.02 x 10²³ particles.
• Treat it like a very large dozen
• 6.02 x 10²³ is called Avogadro's number.

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Representative particles

• The smallest pieces of a substance.
• For an element it is an atom.
• Unless it is diatomic
• For a molecular compound it is a molecule.
• For an ionic compound it is a formula unit.

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Conversion factors

• Used to change units.
• Three questions
• What unit do you want to get rid of?
• Where does it go to cancel out?
• What can you change it into?

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Calculation question

• How many molecules of CO₂ are the in 4.56 moles of CO₂ ?

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Calculation question

• How many moles of water is 5.87 x 10²² molecules?

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Calculation question

• How many atoms of carbon are there in 1.23 moles of C₆H₁₂O₆ ?

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Measuring Moles

• The amu was one twelfth the mass of a carbon 12 atom.
• Since the mole is the number of atoms in 12 grams of carbon-12,
• the decimal number on the periodic table is
• The mass of the average atom in amu
• the mass of 1 mole of those atoms in grams.

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Gram Atomic Mass

• The mass of 1 mole of an element in grams.
• 12.01 grams of carbon has the same number of atoms as 1.01 grams of hydrogen and 55.85 grams of iron.
• We can write this as 12.01 g C = 1 mole
• We can count things by weighing them.

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Examples

• How much would 2.34 moles of carbon weigh?

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Examples

• How many moles of magnesium in 4.61 g of Mg?

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Examples

• How many atoms of lithium in 1.00 g of Li?

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Examples

• How much would 3.45 x 10²² atoms of U weigh?

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What about compounds?

• in 1 mole of H₂O molecules there are two moles of H atoms and 1 mole of O atoms
• To find the mass of one mole of a compound
• determine the moles of the elements they have
• Find out how much they would weigh
• add them up

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What about compounds?

• What is the mass of one mole of CH₄?
• 1 mole of C = 12.01 g
• 4 mole of H x 1.01 g = 4.04g
• 1 mole CH₄ = 12.01 + 4.04 = 16.05g

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Molar Mass

• The mass of 1 mole
• What is the molar mass of Fe₂O₃?
• 2 moles of Fe x 55.85 g = 111.70 g
• 3 moles of O x 16.00 g = 48.00 g
• The GFM = 111.70 g + 48.00 g = 159.70g

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Calculate the molar mass of the following

• C₆H₁₂O₆
• (NH₄)₃PO₄

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Using Molar Mass

Finding moles of compounds

Counting pieces by weighing

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Molar Mass

• The number of grams in 1 mole of atoms, formula units, or molecules.
• We can make conversion factors from these.
• To change grams of a compound to moles of a compound.
• Or moles to grams

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For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles
• for NaOH
• 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g
• 1 mole NaOH = 40.00 g

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For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles
• for NaOH
• 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g
• 1 mole NaOH = 40.00 g

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Gases and the Mole

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Gases

• Many of the chemicals we deal with are gases.
• They are difficult to weigh, so we’ll measure volume
• Need to know how many moles of gas we have.
• Two things affect the volume of a gas
• Temperature and pressure
• Compare at the same temp. and pressure.

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Standard Temperature and Pressure

• Avogadro's Hypothesis - at the same temperature and pressure equal volumes of gas have the same number of particles.
• 0ºC and 1 atmosphere pressure
• Abbreviated atm
• 273 K and 101.3 kPa
• kPa is kiloPascal

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At Standard Temperature and Pressure

• abbreviated STP
• At STP 1 mole of gas occupies 22.4 L
• Called the molar volume
• Used for conversion factors
• Moles to Liter and L to mol

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Examples

• What is the volume of 4.59 mole of CO₂ gas at STP?

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Density of a gas

• D = m /V
• for a gas the units will be g / L
• We can determine the density of any gas at STP if we know its formula.
• To find the density we need the mass and the volume.
• If you assume you have 1 mole than the mass is the molar mass (PT)
• At STP the volume is 22.4 L.

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Examples

• Find the density of CO₂ at STP.

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Quizdom

• Find the density of CH₄ at STP.

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The other way

• Given the density, we can find the molar mass of the gas.
• Again, pretend you have a mole at STP, so V = 22.4 L.
• m = D x V
• m is the mass of 1 mole, since you have 22.4 L of the stuff.
• What is the molar mass of a gas with a density of 1.964 g/L?

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All the things we can change

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PT

Moles

6.02 x 10²³

22.4 L

Count

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Percent Composition

• Like all percents
• Part x 100 % whole
• Find the mass of each component,
• divide by the total mass.

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Example

• Calculate the percent composition of a compound that is 29.0 g of Ag with 4.30 g of S.

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Getting it from the formula

• If we know the formula, assume you have 1 mole.
• Then you know the pieces and the whole.

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Examples

• Calculate the percent composition of C₂H₄?

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Examples

• What is the percent composition of Aluminum carbonate.

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Percent to Mass

• Multiply % by the total mass to find the mass of that component.
• How much aluminum in 450 g of aluminum carbonate?

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Empirical Formula

From percentage to formula

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The Empirical Formula

• The lowest whole number ratio of elements in a compound.
• The molecular formula the actual ratio of elements in a compound.
• The two can be the same.
• CH₂ empirical formula
• C₂H₄ molecular formula
• C₃H₆ molecular formula
• H₂O both

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Finding Empirical Formulas

• Just find the lowest whole number ratio
• C₆H₁₂O₆
• CH₄N₂
• It is not just the ratio of atoms, it is also the ratio of moles of atoms.

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Calculating Empirical Formulas

• Means we can get ratio from percent composition.
• Assume you have a 100 g.
• The percentages become grams.
• Turn grams to moles.
• Find lowest whole number ratio by dividing everything by the smallest moles.

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Example

• Calculate the empirical formula of a compound composed of 38.67 % C, 16.22 % H, and 45.11 %N.
• Assume 100 g so
• 38.67 g C x 1mol C = 3.220 mole C 12.01 gC
• 16.22 g H x 1mol H = 16.1 mole H 1.01 gH
• 45.11 g N x 1mol N = 3.220 mole N 14.01 gN

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Example

• The ratio is 3.220 mol C = 1 mol C 3.220 molN 1 mol N
• The ratio is 16.1 mol H = 5 mol H 3.220 molN 1 mol N
• C₁H₅N₁
• Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula?

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Empirical to molecular

• Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula?
• Since the empirical formula is the lowest ratio the actual molecule would weigh the same or more.
• By a whole number multiple.
• Divide the actual molar mass by the the mass of one mole of the empirical formula.
• You will get a whole number.
• Multiply the empirical formula by this.

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Example

• A compound has an empirical formula of ClCH₂ and a molar mass of 98.96 g/mol. What is its molecular formula?
• A compound has an empirical formula of CH₂O and a molar mass of 180.0 g/mol. What is its molecular formula?

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Percent to molecular

• Take the percent x the molar mass
• This gives you mass in one mole of the compound
• Change this to moles
• You will get whole numbers
• These are the subscripts
• Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. It has a molar mass of 194 g. What is its molecular formula?

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Example

• Ibuprofen is 75.69 % C, 8.80 % H, 15.51 % O, and has a molar mass of about 207 g/mol. What is its molecular formula?

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