Transcript Masterton and Hurley Chapter 3

William L Masterton
Cecile N. Hurley
http://academic.cengage.com/chemistry/masterton
Chapter 3
Mass Relations in Chemistry; Stoichiometry
Edward J. Neth • University of Connecticut
3.1 Atomic Masses
• Atomic mass – (atomic weight) – The atomic mass of
an element indicates how heavy, on average, an
atom of an element is when compared to an atom of
another element
• Atomic mass units – (amu) – the units for atomic
masses on the periodic table
The Carbon-12 Scale
• Mass of one 12C atom = 12 amu (exactly)
• Note that 12C and C-12 mean the same thing
Atomic Masses and Isotopic Abundances:
• Mass spectrometer – a device used to
experimentally determine the atomic mass of an
atom
• Isotopic abundances – the percentage of each
isotope that exists in nature (also, determined using
the mass spec.)
Figure 3.1 – Mass Spectrometer
• A mass spectrometer is used to determine atomic
masses
Figure 3.2 – Mass Spectrum of Cl
• The area under the peak in the mass spectrogram
gives the isotopic abundance
Atomic Mass Calculations:
atomic mass Y  (atomic mass Y1 )
%Y1
%Y2
 (atomic mass Y2 )
100
100
Example 3.1
Masses of Individual Atoms; Avogadro’s Number:
• Avogadro’s Number – The number of atoms that is
equal to the atomic mass of any element
• NA = 6.02 X 1023
• For Example:
• 6.02x1023 H atoms in 1.008 grams of H (atomic
mass of H = 1.008)
Figure 3.3 – One Mole of Several Substances
Example 3.2
3.2 The Mole
• Mole – equal to Avogadro’s Number, equal to
6.02x1023 particles of a substance
Specialized units:
• The correct name for a particle of a substance based
on the type of matter
Atom – the representative particle for an element
example – Fe, S, etc.
ion – the representative particle for a charged particle
example – Na+1, Cl-1, NH4+1,etc.
Molecule – the representative particle for a molecular
compound (made up of non-metals)
example – CO2, CH4, etc.
Formula unit – the representative particle for an ionic
compound (metal and non-metal or polyatomic ion)
example – KCl, MgSO4, etc.
Molar mass:
• Molar mass – (MM) – the mass of 1 mole of a
substance; equal to the atomic mass on the periodic
table
• Round to the tenths place from the Periodic Table to
simplify calculations
Calculating molar mass:
• 1. find the mass of the element on the periodic table
• 2. multiply by the number of atoms of that element in
the formula (distribute parenthesis)
• 3. sum the relative masses of the individual elements
Molar Masses of Some Substances
Practice:
• Example: Determine the molar masses of the
following substances.
Aluminum
CaSO4
(NH4)3P
The Significance of the Mole
• In the laboratory, substances are weighed on
balances, in units of grams
• The mole allows us to relate the number of grams of
a substance to the number of atoms or molecules of
a substance
Mole-Gram Conversions
• Molar mass can be used like any other conversion
factor:
• Molar mass (g) = 1 mole
Example 3.3
3.3 Mass Relations in Chemical Formulas
• Percent composition from formula
• percent composition - the number of grams of
each element in 100 g of the compound
Part x 100 = % composition
whole
2 types of % Comp. problems:
1. Given data as masses of elements, etc.
a. Use the masses of each element and the total mass
of the compound
2 types of % Comp. problems:
2. Given the chemical formula of the compound
a. use the relative molar masses of each element
and the molar mass of the compound
Example 3.4
Subscripts:
1. Represent the atom ratio in a compound
2. Represent the mole ratio in a compound
Diatomic Elements:
• Elements that due to there chemical reactivity exist
only as molecules of 2 atoms in nature.
• 7 diatomic elements:
• Br I N Cl H O F
• H O N Cl Br I F
Simplest Formula from Chemical Analysis
(Empirical Formula):
Simplest formula – (empirical formula) – the simplest
whole number ratio of the atoms in a compound
Calculating the empirical formula:
1. can be determined from masses of the individual elements
or the % composition of the elements in a compound
2. if %’s are given consider the sample to be of 100 grams
and so the %’s become the masses in grams:
•
25.6% = 25.6 g
3. convert the mass of each element to moles
4. divide each number of moles by the smallest number of
moles of all of the answers to #3
5. *If the answers to #4 are whole numbers, these are the
subscripts in the empirical formula.
* If any of the answers to #4 is not a whole number,
convert all answers to a common fraction. Multiply each
fraction by the denominator resulting in a whole number and
these are the subscripts in the empirical formula.
Common Fractions:
0.25
0.33
0.50
0.66
0.75
¼
1/3
½
2/3
¾
Example 3.5 – Simplest Formula from Masses of
Elements
Example 3.6 – Simplest Formula from Mass
Percents
The compound that gives vinegar its sour taste is acetic acid, which contains
the elements carbon, hydrogen, and oxygen. When 5.00g of acetic acid is
analyzed it is found to contain 2.00g of carbon, 0.336g of hydrogen, and
2.66g of oxygen. What is the empirical formula of acetic acid?
Molecular Formula:
• Molecular formula – the true ratio of the atoms in a
compound as it exists naturally; may be the same as
the empirical formula; will relate to the empirical
formula in whole number ratios
Molecular Formula:
Calculating the molecular formula:
1. find the molar mass of the empirical formula
2. divide the molecular molar mass (given in the
question) by the empirical molar mass
MMM
EMM
3. multiply each subscript in the empirical formula by
the answer to #2, these are the subscripts for the
molecular formula
Example 3.7