Download Document

Chemical Compounds
Section 1
Measuring Matter
Section 2
Mass and the Mole
Section 3
Moles of Compounds
Section 4
Empirical and
Molecular Formulas
Section 5
Oxidation States
Section 6
Naming Compounds
Exit
Section1 Measuring Matter
• Explain how a mole is
used to indirectly count
the number of particles of
matter.
molecule: two or more
atoms that covalently
bond together to form a
unit
• Relate the mole to a
common everyday
counting unit.
mole
• Convert between moles
and number of
representative particles.
Avogadro’s number
Chemists use the mole to count atoms,
molecules, ions, and formula units.
Counting Particles
• Chemists need a convenient method for
accurately counting the number of atoms,
molecules, or formula units of a substance.
• The mole is the SI base unit used to measure
the amount of a substance.
• 1 mole is the amount of atoms in 12 g of pure
carbon-12, or 6.02  1023 atoms.
• The number is called Avogadro’s number.
Converting Between Moles and Particles
• Conversion factors must be used.
• Moles to particles
Number of molecules in 3.50 mol of sucrose
Converting Between Moles and Particles (cont.)
• Particles to moles
• Use the inverse of Avogadro’s number as the
conversion factor.
Section1 Assessment
What does the mole measure?
A. mass of a substance
B. amount of a substance
C. volume of a gas
D
C
A
0%
B
D. density of a gas
A. A
B. B
C. C
0%
0%
0%
D. D
Section1 Assessment
What is the conversion factor for
determining the number of moles of a
substance from a known number of
particles?
A
D. 1 mol  6.02  1023 particles
0%
D
C. 1 particle  6.02  1023
C
B.
A. A
B. B
C. C
0%
0%
0%
D. D
B
A.
Section 2 Mass and the Mole
• Relate the mass of an atom conversion factor: a
to the mass of a mole of
ratio of equivalent
atoms.
values used to express
the same quantity in
• Convert between number
different units
of moles and the mass of
an element.
• Convert between number
of moles and number of
atoms of an element.
molar mass
A mole always contains the same number
of particles; however, moles of different
substances have different masses.
The Mass of a Mole
• 1 mol of copper and 1 mol of carbon have
different masses.
• One copper atom has a different mass than 1
carbon atom.
The Mass of a Mole (cont.)
• Molar mass is the mass in grams of one
mole of any pure substance.
• The molar mass of any element is
numerically equivalent to its atomic mass and
has the units g/mol.
Using Molar Mass
• Moles to mass
3.00 moles of copper has a mass of 191 g.
Using Molar Mass (cont.)
• Convert mass to moles with the inverse
molar mass conversion factor.
• Convert moles to atoms with Avogadro’s
number as the conversion factor.
Using Molar Mass (cont.)
• This figure shows the steps to complete
conversions between mass and atoms.
Section2 Assessment
The mass in grams of 1 mol of any pure
substance is:
A. molar mass
B. Avogadro’s number
D
A
0%
C
D. 1 g/mol
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. atomic mass
Section 2 Assessment
Molar mass is used to convert what?
A. mass to moles
B. moles to mass
C. atomic weight
D
C
A
0%
B
D. particles
A. A
B. B
C. C
0%
0%
0%
D. D
Section 3 Moles of Compounds
• Recognize the mole relationships shown by a
chemical formula.
• Calculate the molar mass of a compound.
• Convert between the number of moles and mass of
a compound.
• Apply conversion factors to determine the number of
atoms or ions in a known mass of a compound.
representative particle: an atom, molecule, formula
unit, or ion
Chemical Formulas and the Mole
• Chemical formulas indicate the numbers
and types of atoms contained in one unit of
the compound.
• One mole of CCl2F2 contains one mole of C
atoms, two moles of Cl atoms, and two moles
of F atoms.
The Molar Mass of Compounds
• The molar mass of a compound equals the
molar mass of each element, multiplied by
the moles of that element in the chemical
formula, added together.
• The molar mass of a compound
demonstrates the law of conservation of
mass.
Converting Moles of a Compound to Mass
• For elements, the conversion factor is the
molar mass of the compound.
• The procedure is the same for compounds,
except that you must first calculate the molar
mass of the compound.
Converting the Mass of a Compound to Moles
• The conversion factor is the inverse of the
molar mass of the compound.
Converting the Mass of a Compound to
Number of Particles
• Convert mass to moles of compound with
the inverse of molar mass.
• Convert moles to particles with Avogadro’s
number.
Converting the Mass of a Compound to
Number of Particles (cont.)
• This figure summarizes the conversions
between mass, moles, and particles.
Section 3 Assessment
How many moles of OH— ions are in 2.50
moles of Ca(OH)2?
A. 2.00
B. 2.50
D
A
0%
C
D. 5.00
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. 4.00
Section 3 Assessment
How many particles of Mg are in 10 moles
of MgBr2?
A. 6.02  1023
B. 6.02  1024
D
A
0%
C
D. 1.20  1025
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. 1.20  1024
Section 4 Empirical and Molecular Formulas
• Explain what is meant by
the percent composition
of a compound.
• Determine the
empirical and molecular
formulas for a
compound from mass
percent and actual
mass data.
percent by mass: the
ratio of the mass of each
element to the total
mass of the compound
expressed as a percent
percent composition
empirical formula
molecular formula
A molecular formula of a compound is
a whole-number multiple of its
empirical formula.
Percent Composition
• The percent by mass of any element in a
compound can be found by dividing the
mass of the element by the mass of the
compound and multiplying by 100.
Percent Composition (cont.)
• The percent by mass of each element in a
compound is the percent composition of
a compound.
• Percent composition of a compound can also
be determined from its chemical formula.
Empirical Formula
• The empirical formula for a compound is
the smallest whole-number mole ratio of
the elements.
• You can calculate the empirical formula from
percent by mass by assuming you have
100.00 g of the compound. Then, convert the
mass of each element to moles.
• The empirical formula may or may not be the
same as the molecular formula.
Molecular formula of hydrogen peroxide = H2O2
Empirical formula of hydrogen peroxide = HO
Molecular Formula
• The molecular formula specifies the
actual number of atoms of each element in
one molecule or formula unit of the
substance.
• Molecular formula is always a whole-number
multiple of the empirical formula.
Molecular Formula (cont.)
Chemical Composition
Halothane
C2HBrClF3
Mole ratio
nC/nhalothane
Mass ratio
mC/mhalothane
M(C2HBrClF3) = 2MC + MH + MBr + MCl + 3MF
= (2  12.01) + 1.01 + 79.90 + 35.45 + (3  19.00)
= 197.38 g/mol
Example
Calculating the Mass Percent Composition of a Compound
Calculate the molecular mass
M(C2HBrClF3) = 197.38 g/mol
For one mole of compound, formulate the mass
ratio and convert to percent:
(2 12.01) g
%C 
100%  12.17%
197.38 g
Example
(2  12.01) g
%C 
100%  12.17%
197.38 g
1.01g
%H 
100%  0.51%
197.38 g
79.90 g
% Br 
100%  40.48%
197.38 g
35.45 g
%Cl 
 100%  17.96%
197.38 g
(3 19.00) g
%F 
100%  28.88%
197.38 g
Empirical formula
5 Step approach:
1.
2.
3.
4.
5.
Choose an arbitrary sample size (100g).
Convert masses to amounts in moles.
Write a formula.
Convert formula to small whole numbers.
Multiply all subscripts by a small whole number
to make the subscripts integral.
Example
Determining the Empirical and Molecular Formulas of a
Compound from Its Mass Percent Composition.
Dibutyl succinate is an insect repellent used against household
ants and roaches. Its composition is 62.58% C, 9.63% H and
27.79% O. Its experimentally determined molecular mass is
230 u. What are the empirical and molecular formulas of
dibutyl succinate?
Step 1: Determine the mass of each element in a 100g sample.
C 62.58 g
H 9.63 g O 27.79 g
Example
Step 2: Convert masses to amounts in moles.
1 mol C
 5.210 mol C
12.011 g C
1 mol H
nH  9.63 g H 
 9.55 mol H
1.008 g H
1 mol O
nO  27.79 g O 
 1.737 mol O
15.999 g O
nC  62.58 g C 
Step 3: Write a tentative formula.
C5.21H9.55O1.74
Step 4: Convert to small whole numbers.
C2.99H5.49O
Example
Step 5: Convert to a small whole number ratio.
Multiply 2 to get C5.98H10.98O2
The empirical formula is C6H11O2
Step 6: Determine the molecular formula.
Empirical formula mass is 115 u.
Molecular formula mass is 230 u.
The molecular formula is C12H22O4
Oxidation States
Metals tend to
lose electrons.
Non-metals tend
to gain electrons.
Na  Na+ + e-
Cl + e-  Cl-
Reducing agents
Oxidizing agents
We use the Oxidation State to keep track of the number of
electrons that have been gained or lost by an element.
Rules for Oxidation States
1.
The oxidation state (OS) of an individual atom in a free
element is 0.
2.
The total of the OS in all atoms in:
i. Neutral species is 0.
ii. Ionic species is equal to the charge on the ion.
3.
In their compounds, the alkali metals and the alkaline
earths have OS of +1 and +2 respectively.
4.
In compounds the OS of fluorine is always –1
Rules for Oxidation States
5.
In compounds, the OS of hydrogen is usually +1
6.
In compounds, the OS of oxygen is usually –2.
7.
In binary (two-element) compounds with metals:
i. Halogens have OS of –1,
ii. Group 16 have OS of –2 and
iii. Group 15 have OS of –3.
Example
Assigning Oxidation States.
What is the oxidation state of the underlined element in each
of the following? a) P4; b) Al2O3; c) MnO4-; d) NaH
a) P4 is an element. P OS = 0
b) Al2O3: O is –2. O3 is –6. Since (+6)/2=(+3), Al OS = +3.
c) MnO4-: net OS = -1, O4 is –8. Mn OS = +7.
d) NaH: net OS = 0, rule 3 beats rule 5, Na OS = +1 and
H OS = -1.
Naming Compounds
Trivial names are used for common compounds.
A systematic method of naming compounds is
known as a system of nomenclature.
Inorganic compounds
Inorganic Nomenclature
Binary Compounds of Metals and Nonmetals
NaCl
=
electrically neutral
sodium chloride
name is unchanged
“ide” ending
MgI2
=
magnesium iodide
Al2O3
=
aluminum oxide
Na2S
=
sodium sulfide
Binary Compounds of Two Non-metals
Molecular compounds
usually write the positive OS element first.
HCl hydrogen chloride
Some pairs form more than one compound
mono 1
penta
5
di
2
hexa
6
tri
3
hepta
7
tetra
4
octa
8
Binary Acids
Acids produce H+ when dissolved in water.
They are compounds that ionize in water.
Emphasize the fact that a molecule is an acid by altering the name.
HCl
hydrogen chloride
hydrochloric acid
HF
hydrogen fluoride
hydrofluoric acid
Polyatomic Ions
Polyatomic ions are very common.
Table 3.3 gives a list of some of them. Here are a few:
ammonium ion
NH4+
acetate ion
C2H3O2-
carbonate ion
CO32-
hydrogen carbonate
HCO3-
hypochlorite
ClO-
phosphate
PO43-
chlorite
ClO2-
hydrogen phosphate
HPO42-
chlorate
ClO3-
sulfate
SO42-
perchlorate
ClO4-
hydrogensulfate
HSO4-
•Slide 49 of 37
•General Chemistry: Chapter 3
•Prentice-Hall © 2002
How many moles of hydrogen atoms are
in one mole of H2O2?
A. 1
B. 2
D
A
0%
C
D. 0.5
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. 3
How many moles of Al are in 2.0 mol of
Al2Br3?
A. 2
B. 4
D
A
0%
C
D. 1
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. 6
How many atoms of hydrogen are in
3.5 mol of H2S?
A. 7.0  1023
B. 2.1  1023
D
A
0%
C
D. 4.2  1024
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. 6.0  1023
Inorganic Nomenclature
• Write the name of the cation.
• If the anion is an element, change its
ending to -ide; if the anion is a
polyatomic ion, simply write the name of
the polyatomic ion.
• If the cation can have more than one
possible charge, write the charge as a
Roman numeral in parentheses.
Atoms,
Molecules,
and Ions
•© 2009, Prentice-Hall,
Patterns in Oxyanion Nomenclature
• When there are two oxyanions involving
the same element:
– The one with fewer oxygens ends in -ite.
• NO2− : nitrite; SO32− : sulfite
– The one with more oxygens ends in -ate.
• NO3− : nitrate; SO42− : sulfate
Atoms,
Molecules,
and Ions
•© 2009, Prentice-Hall,
Patterns in Oxyanion
Nomenclature
• The one with the second fewest oxygens ends in -ite.
– ClO2− : chlorite
• The one with the second most oxygens ends in -ate.
– ClO3− : chlorate
Atoms,
Molecules,
and Ions
•© 2009, Prentice-Hall,
Patterns in Oxyanion Nomenclature
• The one with the fewest oxygens has the prefix hypoand ends in -ite.
– ClO− : hypochlorite
• The one with the most oxygens has the prefix per- and
ends in -ate.
– ClO4− : perchlorate
Atoms,
Molecules,
and Ions
•© 2009, Prentice-Hall,
Acid Nomenclature
• If the anion in the acid
ends in -ide, change
the ending to -ic acid
and add the prefix
hydro- .
– HCl: hydrochloric acid
– HBr: hydrobromic acid
– HI: hydroiodic acid
Atoms,
Molecules,
and Ions
•© 2009, Prentice-Hall,
Acid Nomenclature
• If the anion in the acid
ends in -ite, change
the ending to -ous
acid.
– HClO: hypochlorous
acid
– HClO2: chlorous acid
Atoms,
Molecules,
and Ions
•© 2009, Prentice-Hall,
Acid Nomenclature
• If the anion in the acid
ends in -ate, change
the ending to -ic acid.
– HClO3: chloric acid
– HClO4: perchloric acid
Atoms,
Molecules,
and Ions
•© 2009, Prentice-Hall,
Nomenclature of Binary
Compounds
• The less electronegative
atom is usually listed first.
• A prefix is used to denote
the number of atoms of
each element in the
compound (mono- is not
used on the first element
listed, however) .
Atoms,
Molecules,
and Ions
•© 2009, Prentice-Hall,
Nomenclature of Binary
Compounds
• The ending on the more
electronegative element
is changed to -ide.
– CO2: carbon dioxide
– CCl4: carbon tetrachloride
Atoms,
Molecules,
and Ions
•© 2009, Prentice-Hall,
Nomenclature of Binary
Compounds
• If the prefix ends with a
or o and the name of the
element begins with a
vowel, the two
successive vowels are
often elided into one.
N2O5: dinitrogen pentoxide
Atoms,
Molecules,
and Ions
•© 2009, Prentice-Hall,