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Chapter 7 Chemical Quantities

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Calculate the percent composition of a substance from its chemical formula or experimental data. ... The percentages become grams. Convert grams to moles. ... – PowerPoint PPT presentation

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Title: Chapter 7 Chemical Quantities


1
Chapter 7Chemical Quantities
  • Pioneer High School
  • Mr. Gonzalez

2
Section 7.1The Mole A Measurement of Matter
  • OBJECTIVES
  • Describe how Avogadros number is related to a
    mole of any substance.

3
Section 7.1The Mole A Measurement of Matter
  • OBJECTIVES
  • Calculate the mass of a mole of any substance.

4
What is a Mole?
  • 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.

5
Moles (abbreviated mol)
  • 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 Avogadros number.

6
Representative particles
  • The smallest pieces of a substance.
  • For a molecular compound it is the molecule.
  • For an ionic compound it is the formula unit
    (ions).
  • For an element it is the atom.
  • Remember the 7 diatomic elements (made of
    molecules)

7
Types of questions
  • How many oxygen atoms in the following?
  • CaCO3
  • Al2(SO4)3
  • How many ions in the following?
  • CaCl2
  • NaOH
  • Al2(SO4)3

8
Types of questions
  • How many molecules of CO2 are there in 4.56 moles
    of CO2 ?
  • How many moles of water is 5.87 x 1022
    molecules?
  • How many atoms of carbon are there in 1.23 moles
    of C6H12O6 ?
  • How many moles is 7.78 x 1024 formula units of
    MgCl2?

9
Measuring Moles
  • Remember relative atomic mass?
  • 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 also
    the mass of 1 mole of those atoms in grams.

10
Gram Atomic Mass (gam)
  • Equals the mass of 1 mole of an element in grams
  • 12.01 grams of C has the same number of pieces as
    1.008 grams of H and 55.85 grams of iron.
  • We can write this as 12.01 g C 1 mole C
  • We can count things by weighing them.

11
Examples
  • How much would 2.34 moles of carbon weigh?
  • How many moles of magnesium is 24.31 g of Mg?
  • How many atoms of lithium is 1.00 g of Li?
  • How much would 3.45 x 1022 atoms of U weigh?

12
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

13
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
  • The Gram Molecular Mass (gmm) of CH4 is 16.05g
  • this is the mass of one mole of a molecular
    compound.

14
Gram Formula Mass (gfm)
  • The mass of one mole of an ionic compound.
  • Calculated the same way as gmm.
  • What is the GFM 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 159.70 g

15
Section 7.2Mole-Mass and Mole-Volume
Relationships
  • OBJECTIVES
  • Use the molar mass to convert between mass and
    moles of a substance.

16
Section 7.2Mole-Mass and Mole-Volume
Relationships
  • OBJECTIVES
  • Use the mole to convert among measurements of
    mass, volume, and number of particles.

17
Molar Mass
  • Molar mass is the generic term for the mass of
    one mole of any substance (in grams)
  • The same as 1) gram molecular mass, 2) gram
    formula mass, and 3) gram atomic mass- just a
    much broader term.

18
Examples
  • Calculate the molar mass of the following and
    tell what type it is
  • Na2S
  • N2O4
  • C
  • Ca(NO3)2
  • C6H12O6
  • (NH4)3PO4

19
Molar Mass
  • The number of grams of 1 mole of atoms, ions, or
    molecules.
  • We can make conversion factors from these.
  • To change grams of a compound to moles of a
    compound.

20
For example
  • How many moles is 5.69 g of NaOH?

21
For example
  • How many moles is 5.69 g of NaOH?

22
For example
  • How many moles is 5.69 g of NaOH?
  • need to change grams to moles

23
For example
  • How many moles is 5.69 g of NaOH?
  • need to change grams to moles
  • for NaOH

24
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

25
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

26
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

27
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

28
Examples
  • How many moles is 4.56 g of CO2?
  • How many grams is 9.87 moles of H2O?
  • How many molecules is 6.8 g of CH4?
  • 49 molecules of C6H12O6 weighs how much?

29
Gases
  • Many of the chemicals we deal with are gases.
  • They are difficult to weigh.
  • Need to know how many moles of gas we have.
  • Two things effect the volume of a gas
  • Temperature and pressure
  • We need to compare them at the same temperature
    and pressure.

30
Standard Temperature and Pressure
  • 0ºC and 1 atm pressure
  • abbreviated STP
  • At STP 1 mole of gas occupies 22.4 L
  • Called the molar volume
  • 1 mole 22.4 L of any gas at STP

31
Examples
  • What is the volume of 4.59 mole of CO2 gas at
    STP?
  • How many moles is 5.67 L of O2 at STP?
  • What is the volume of 8.8 g of CH4 gas at STP?

32
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, then the mass is
    the molar mass (from PT)
  • At STP the volume is 22.4 L.

33
Examples
  • Find the density of CO2 at STP.
  • Find the density of CH4 at STP.

34
The other way
  • Given the density, we can find the molar mass of
    the gas.
  • Again, pretend you have 1 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?
  • 2.86 g/L?

35
Summary
  • These four items are all equal
  • a) 1 mole
  • b) molar mass (in grams)
  • c) 6.02 x 1023 representative particles
  • d) 22.4 L at STP
  • Thus, we can make conversion factors from them.

36
Section 7.3Percent Composition and Chemical
Formulas
  • OBJECTIVES
  • Calculate the percent composition of a substance
    from its chemical formula or experimental data.

37
Section 7.3Percent Composition and Chemical
Formulas
  • OBJECTIVES
  • Derive the empirical formula and the molecular
    formula of a compound from experimental data.

38
Calculating Percent Composition of a Compound
  • Like all percent problems
  • Part whole
  • Find the mass of each component,
  • then divide by the total mass.

x 100
39
Example
  • Calculate the percent composition of a compound
    that is 29.0 g of Ag with 4.30 g of S.

40
Getting it from the formula
  • If we know the formula, assume you have 1 mole.
  • Then you know the mass of the pieces and the
    whole.

41
Examples
  • Calculate the percent composittion of C2H4?
  • How about Aluminum carbonate?
  • Sample Problem 7-11, p.191
  • We can also use the percent as a conversion
    factor
  • Sample Problem 7-12, p.191

42
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 is an empirical formula
  • C2H4 is a molecular formula
  • C3H6 is a molecular formula
  • H2O is both empirical molecular

43
Calculating Empirical
  • Just find the lowest whole number ratio
  • C6H12O6
  • CH4N
  • It is not just the ratio of atoms, it is also the
    ratio of moles of atoms.
  • In 1 mole of CO2 there is 1 mole of carbon and 2
    moles of oxygen.
  • In one molecule of CO2 there is 1 atom of C and 2
    atoms of O.

44
Calculating Empirical
  • We can get a ratio from the percent composition.
  • Assume you have a 100 g.
  • The percentages become grams.
  • Convert grams to moles.
  • Find lowest whole number ratio by dividing by the
    smallest.

45
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 16.09 mole H 1.01
    gH
  • 45.11 g N x 1mol N 3.219 mole N 14.01
    gN

46
Example
  • The ratio is 3.220 mol C 1 mol C
    3.219 molN 1 mol N
  • The ratio is 16.09 mol H 5 mol H
    3.219 molN 1 mol N
  • C1H5N1
  • A compound is 43.64 P and 56.36 O. What is
    the empirical formula?
  • Caffeine is 49.48 C, 5.15 H, 28.87 N and
    16.49 O. What is its empirical formula?

47
Empirical to molecular
  • Since the empirical formula is the lowest ratio,
    the actual molecule would weigh more.
  • By a whole number multiple.
  • Divide the actual molar mass by the empirical
    formula mass.
  • Caffeine has a molar mass of 194 g. what is its
    molecular formula?

48
Example
  • A compound is known to be composed of 71.65 Cl,
    24.27 C and 4.07 H. Its molar mass is known
    (from gas density) to be 98.96 g. What is its
    molecular formula?
  • Sample Problem 7-14, p.194
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