cmc chapter 10

Chapter Menu

The Mole

Section 10.1 Measuring Matter

Section 10.2 Mass and the Mole

Section 10.3 Moles of Compounds

Section 10.4 Empirical and Molecular Formulas

Section 10.5 Formulas of Hydrates

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Section 10-1


• 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

mole

Avogadro’s number

• Relate the mole to a common everyday counting unit.

• Convert between moles and number of representative particles.

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 1023 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.

A. A

B. B

C. C

D. D

Section 10-1

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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. A

B. B

C. C

D. D

Section 10-1

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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 1023

D. 1 mol 6.02 1023 particles

End of Section 10-1

Section 10-2


• Relate the mass of an atom to the mass of a mole of atoms.

conversion factor: a ratio of equivalent values used to express the same quantity in different units

molar mass

• Convert between number of moles and the mass of an element.

• Convert between number of moles and number of atoms of an element.

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.

A. A

B. B

C. C

D. D

Section 10-2

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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. A

B. B

C. C

D. D

Section 10-2


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Molar mass is used to convert what?

A. mass to moles

B. moles to mass

C. atomic weight

D. particles

End of Section 10-2

Section 10-3


• Recognize the mole relationships shown by a chemical formula.

representative particle: an atom, molecule, formula unit, or ion

• 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.

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 CCl2F2 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.

A. A

B. B

C. C

D. D

Section 10-3

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How many moles of OH— ions are in 2.50 moles of Ca(OH)2?

A. 2.00

B. 2.50

C. 4.00

D. 5.00

A. A

B. B

C. C

D. D

Section 10-3


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How many particles of Mg are in 10 moles of MgBr2?

A. 6.02 1023

B. 6.02 1024

C. 1.20 1024

D. 1.20 1025

End of Section 10-3

Section 10-4


• Explain what is meant by the percent composition of a compound.

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

• Determine the empirical and molecular formulas for a compound from mass percent and actual mass data.

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 percent composition of a compound.

• 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 = H2O2

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.)

A. A

B. B

C. C

D. D

Section 10-4

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What is the empirical formula for the compound C6H12O6?

A. CHO

B. C2H3O2

C. CH2O

D. CH3O

A. A

B. B

C. C

D. D

Section 10-4


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Which is the empirical formula for hydrogen peroxide?

A. H2O2

B. H2O

C. HO

D. none of the above

End of Section 10-4

Section 10-5


• Explain what a hydrate is and relate the name of the hydrate to its composition.

crystal lattice: a three-dimensional geometric arrangement of particles

hydrate

• Determine the formula of a hydrate from laboratory data.

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 = Na2CO3 • 10H2O

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.

A. A

B. B

C. C

D. D

Section 10-5

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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. A

B. B

C. C

D. D

Section 10-5


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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

End of Section 10-5

Resources Menu

Chemistry Online

Study Guide

Chapter Assessment

Standardized Test Practice

Image Bank

Concepts in Motion

http://www.glencoe.com/

Study Guide 1


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 2


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 3


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 4


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 5


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.

A. A

B. B

C. C

D. D

Chapter Assessment 1

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What does Avogadro’s number represent?

A. the number of atoms in 1 mol of an element

B. the number of molecules in 1 mol of a compound

C. the number of Na+ ions in 1 mol ofNaCl (aq)

D. all of the above

A. A

B. B

C. C

D. D


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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. A

B. B

C. C

D. D


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How many moles of hydrogen atoms are in one mole of H2O2?

A. 1

B. 2

C. 3

D. 0.5

A. A

B. B

C. C

D. D


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What is the empirical formula of Al2Br3?

A. AlBr

B. AlBr3

C. Al2Br

D. Al2Br3

A. A

B. B

C. C

D. D


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What is an ionic solid with trapped water molecules called?

A. aqueous solution

B. anhydrous compound

C. hydrate

D. solute

A. A

B. B

C. C

D. D

STP 1

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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. A

B. B

C. C

D. D

STP 2

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How many moles of Al are in 2.0 mol of Al2Br3?

A. 2

B. 4

C. 6

D. 1

A. A

B. B

C. C

D. D

STP 3

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How many water molecules are associated with 3.0 mol of CoCl2 • 6H2O?

A. 18

B. 1.1 1025

C. 3.6 1024

D. 1.8 1024

A. A

B. B

C. C

D. D

STP 4

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How many atoms of hydrogen are in 3.5 mol of H2S?

A. 7.0 1023

B. 2.1 1023

C. 6.0 1023

D. 4.2 1024

A. A

B. B

C. C

D. D

STP 5

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Which is not the correct formula for an ionic compound?

A. CO2

B. NaCl

C. Na2SO4

D. LiBr2

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CIM

Figure 10.6 Molar Mass

Table 10.1 Formulas of Hydrates

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cmc chapter 10

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