Transcript CMC Chapter 10
The Mole
Measuring Matter
Mass and the Mole
Moles of Compounds
Empirical and Molecular Formulas
Formulas of Hydrates
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Section 10.1 Measuring Matter
• • •
Explain
how a mole is used to indirectly count the number of particles of matter.
Relate
the mole to a common everyday counting unit.
Convert
between moles and number of representative particles.
molecule:
atoms that covalently bond together to form a unit mole two or more Avogadro’s number
Chemists use the mole to count atoms, molecules, ions, and formula units.
Section 10-1
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 10 23 atoms.
• The number is called
Avogadro’s number
.
Section 10-1
Converting Between Moles and Particles
• Conversion factors must be used.
• Moles to particles Number of molecules in 3.50 mol of sucrose Section 10-1
Converting Between Moles and Particles (cont.)
• Particles to moles • Use the inverse of Avogadro’s number as the conversion factor.
Section 10-1
Section 10.1 Assessment What does the mole measure?
A.
mass of a substance
B.
amount of a substance
C.
volume of a gas
D.
density of a gas
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Section 10-1
Section 10.1 Assessment What is the conversion factor for determining the number of moles of a substance from a known number of particles? A.
B.
C.
1 particle 6.02 10 23
D.
1 mol 6.02 10 23 particles
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Section 10-1
End of Section 10-1
Section 10.2 Mass and the Mole
• • •
Relate
the mass of an atom to the mass of a mole of atoms.
Convert
between number of moles and the mass of an element.
Convert
between number of moles and number of atoms of an element.
conversion factor:
ratio of equivalent different units a values used to express the same quantity in molar mass
A mole always contains the same number of particles; however, moles of different substances have different masses.
Section 10-2
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.
Section 10-2
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.
Section 10-2
Using Molar Mass
• Moles to mass 3.00 moles of copper has a mass of 191 g.
Section 10-2
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.
Section 10-2
Using Molar Mass (cont.)
• This figure shows the steps to complete conversions between mass and atoms.
Section 10-2
Section 10.2 Assessment The mass in grams of 1 mol of any pure substance is: A.
molar mass
B.
Avogadro’s number
C.
atomic mass
D.
1 g/mol
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Section 10-2
Section 10.2 Assessment Molar mass is used to convert what?
A.
mass to moles
B.
moles to mass
C.
atomic weight
D.
particles
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Section 10-2
End of Section 10-2
Section 10.3 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:
unit, or ion an atom, molecule, formula Section 10-3
Section 10.3 Moles of Compounds (cont.) The molar mass of a compound can be calculated from its chemical formula and can be used to convert from mass to moles of that compound.
Section 10-3
Chemical Formulas and the Mole
• Chemical formulas indicate the numbers and types of atoms contained in one unit of the compound.
• One mole of CCl 2 F 2 contains one mole of C atoms, two moles of Cl atoms, and two moles of F atoms.
Section 10-3
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.
Section 10-3
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.
Section 10-3
Converting the Mass of a Compound to Moles
• The conversion factor is the inverse of the molar mass of the compound.
Section 10-3
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.
Section 10-3
Converting the Mass of a Compound to Number of Particles (cont.)
• This figure summarizes the conversions between mass, moles, and particles.
Section 10-3
Section 10.3 Assessment How many moles of OH — moles of Ca(OH) 2 ?
ions are in 2.50 A.
2.00
B.
2.50
C.
4.00
D.
5.00
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Section 10-3
Section 10.3 Assessment How many particles of Mg are in 10 moles of MgBr 2 ?
A.
6.02 10 23
B.
C.
D.
6.02 10 24 1.20 10 24 1.20 10 25 A. A B. B
A 0% B 0%
C. C
0%
D. D
C 0% D
Section 10-3
End of Section 10-3
Section 10.4 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.
Section 10-4
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.
Section 10-4
Percent Composition (cont.)
• The percent by mass of each element in a compound is the a compound.
percent composition
of • Percent composition of a compound can also be determined from its chemical formula.
Section 10-4
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 = H 2 O 2 Empirical formula of hydrogen peroxide = HO Section 10-4
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.
Section 10-4
Molecular Formula (cont.)
Section 10-4
Section 10.4 Assessment What is the empirical formula for the compound C 6 H 12 O 6 ?
A.
CHO
B.
C 2 H 3 O 2
C.
CH 2 O
D.
CH 3 O
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Section 10-4
Section 10.4 Assessment Which is the empirical formula for hydrogen peroxide?
A.
H 2 O 2
B.
H 2 O
C.
HO
D.
none of the above
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Section 10-4
End of Section 10-4
Section 10.5 Formulas of Hydrates
•
Explain
what a hydrate is and relate the name of the hydrate to its composition.
•
Determine
the formula of a hydrate from laboratory data.
crystal lattice:
a three dimensional geometric arrangement of particles hydrate
Hydrates are solid ionic compounds in which water molecules are trapped.
Section 10-5
Naming Hydrates
• A
hydrate
is a compound that has a specific number of water molecules bound to its atoms.
• The number of water molecules associated with each formula unit of the compound is written following a dot.
• Sodium carbonate decahydrate = Na 2 CO 3 • 10H 2 O Section 10-5
Naming Hydrates (cont.)
Section 10-5
Analyzing a Hydrate
• When heated, water molecules are released from a hydrate leaving an anhydrous compound.
• To determine the formula of a hydrate, find the number of moles of water associated with 1 mole of hydrate.
Section 10-5
Analyzing a Hydrate (cont.)
• Weigh hydrate.
• Heat to drive off the water.
• Weigh the anhydrous compound.
• Subtract and convert the difference to moles.
• The ratio of moles of water to moles of anhydrous compound is the coefficient for water in the hydrate.
Section 10-5
Use of Hydrates
• Anhydrous forms of hydrates are often used to absorb water, particularly during shipment of electronic and optical equipment.
• In chemistry labs, anhydrous forms of hydrates are used to remove moisture from the air and keep other substances dry.
Section 10-5
Section 10.5 Assessment Heating a hydrate causes what to happen?
A.
Water is driven from the hydrate.
B.
The hydrate melts.
C.
The hydrate conducts electricity.
D.
There is no change in the hydrate.
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Section 10-5
Section 10.5 Assessment A hydrate that has been heated and the water driven off is called: A.
dehydrated compound
B.
antihydrated compound
C.
anhydrous compound
D.
hydrous compound
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Section 10-5
End of Section 10-5
Chemistry Online Study Guide Chapter Assessment Standardized Test Practice Image Bank Concepts in Motion Resources Menu
Section 10.1 Measuring Matter Key Concepts
• The mole is a unit used to count particles of matter indirectly. One mole of a pure substance contains Avogadro’s number of particles.
• Representative particles include atoms, ions, molecules, formula units, electrons, and other similar particles.
• One mole of carbon-12 atoms has a mass of exactly 12 g.
• Conversion factors written from Avogadro’s relationship can be used to convert between moles and number of representative particles.
Study Guide 1
Section 10.2 Mass and the Mole Key Concepts
• The mass in grams of 1 mol of any pure substance is called its molar mass.
• The molar mass of an element is numerically equal to its atomic mass.
• The molar mass of any substance is the mass in grams of Avogadro’s number of representative particles of the substance.
• Molar mass is used to convert from moles to mass. The inverse of molar mass is used to convert from mass to moles.
Study Guide 2
Section 10.3 Moles of Compounds Key Concepts
• Subscripts in a chemical formula indicate how many moles of each element are present in 1 mol of the compound.
• The molar mass of a compound is calculated from the molar masses of all of the elements in the compound. • Conversion factors based on a compound’s molar mass are used to convert between moles and mass of a compound.
Study Guide 3
Section 10.4 Empirical and Molecular Formulas Key Concepts
• The percent by mass of an element in a compound gives the percentage of the compound’s total mass due to that element.
• The subscripts in an empirical formula give the smallest whole-number ratio of moles of elements in the compound.
• The molecular formula gives the actual number of atoms of each element in a molecule or formula unit of a substance. • The molecular formula is a whole-number multiple of the empirical formula.
Study Guide 4
Section 10.5 Formulas of Hydrates Key Concepts
• The formula of a hydrate consists of the formula of the ionic compound and the number of water molecules associated with one formula unit.
• The name of a hydrate consists of the compound name and the word hydrate with a prefix indicating the number of water molecules in 1 mol of the compound. • Anhydrous compounds are formed when hydrates are heated.
Study Guide 5
What does Avogadro’s number represent? A.
the number of atoms in 1 mol of an element
B.
C.
D.
the number of molecules in 1 mol of a compound the number of Na + NaCl (aq) ions in 1 mol of all of the above
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Chapter Assessment 1
The molar mass of an element is numerically equivalent to what?
A.
1 amu
B.
1 mole
C.
its atomic mass
D.
its atomic number
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Chapter Assessment 2
How many moles of hydrogen atoms are in one mole of H 2 O 2 ?
A.
1
B.
2
C.
3
D.
0.5
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Chapter Assessment 3
What is the empirical formula of Al 2 Br 3 ?
A.
AlBr
B.
AlBr 3
C.
Al 2 Br
D.
Al 2 Br 3
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Chapter Assessment 4
What is an ionic solid with trapped water molecules called?
A.
aqueous solution
B.
anhydrous compound
C.
hydrate
D.
solute
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
Chapter Assessment 5
Two substances have the same percent by mass composition, but very different properties. They must have the same ____.
A.
density
B.
empirical formula
C.
molecular formula
D.
molar mass
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
STP 1
How many moles of Al are in 2.0 mol of Al 2 Br 3 ?
A.
2
B.
4
C.
6
D.
1
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
STP 2
How many water molecules are associated with 3.0 mol of CoCl 2 • 6H 2 O?
A.
B.
C.
D.
18 1.1 10 25 3.6 10 24 1.8 10 24
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
STP 3
How many atoms of hydrogen are in 3.5 mol of H 2 S?
A.
7.0 10 23
B.
C.
D.
2.1 10 23 6.0 10 23 4.2 10 24
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
STP 4
Which is not the correct formula for an ionic compound?
A.
CO 2
B.
NaCl
C.
Na 2 SO 4
D.
LiBr 2
A 0%
A. A B. B
B 0%
C. C
0%
D. D
C 0% D
STP 5
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Figure 10.6 Molar Mass Table 10.1 Formulas of Hydrates CIM
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