chapter 10: the mole chemistry matter and change

Chapter 10: The MoleChapter 10: The Mole

CHEMISTRY Matter and Change

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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Table Of ContentsCHAPTER

10

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• 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.

SECTION10.1

Measuring Matter

Counting Particles• Chemists need a convenient method for accurately

counting the number of atoms, molecules, or formula units of a substance.

• The mole is *

• 1 mole is * , which is any kind of particle – an atom, a molecule, a formula unit, an electron, an ion, etc.

• The number is called Avogadro’s number.

SECTION10.1

Measuring Matter

Converting Between Moles and Particles• Conversion factors must be used.

• Moles to particles

Number of molecules in 3.50 mol of sucrose

SECTION10.1

Measuring Matter

Converting Between Moles and Particles (cont.)

• Particles to moles

• Use the inverse of Avogadro’s number as the conversion factor.

SECTION10.1

Measuring Matter

• 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.

SECTION10.2

Mass and the Mole

The Mass of a Mole

• 1 mol of copper (6.02 x 1023 atoms of copper) and 1 mol of carbon (6.02 x 1023 atoms of carbon) have different masses.

• One copper atom has a different mass than 1 carbon atom.

SECTION10.2

Mass and the Mole

The Mass of a Mole (cont.)

• Molar mass is *.

• The molar mass of any element is numerically equivalent to its atomic mass and has the units g/mol.

SECTION10.2

Mass and the Mole

Using Molar Mass

• Moles to mass

3.00 moles of copper has a mass of 191 g.

SECTION10.2

Mass and the Mole

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.

SECTION10.2

Mass and the Mole

Using Molar Mass (cont.)

• This figure shows the steps to complete conversions between mass and atoms.

SECTION10.2

Mass and the Mole

• 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.

SECTION10.3

Moles of Compounds

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.

SECTION10.3

Moles of Compounds

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.

SECTION10.3

Moles of Compounds

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.

SECTION10.3

Moles of Compounds

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.

SECTION10.3

Moles of Compounds

Converting the Mass of a Compound to Moles• The conversion factor is the inverse of the

molar mass of the compound.

SECTION10.3

Moles of Compounds

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.

SECTION10.3

Moles of Compounds

Converting the Mass of a Compound to Number of Particles (cont.)

• This figure summarizes the conversions between mass, moles, and particles.

SECTION10.3

Moles of Compounds

• 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.

SECTION10.4

Empirical and Molecular Formulas

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.

SECTION10.4

Empirical and Molecular Formulas

Percent Composition (cont.)

• * is the percent composition of a compound.

• Percent composition of a compound can also be determined from its chemical formula.

SECTION10.4

Empirical and Molecular Formulas

Empirical Formula• The empirical formula for a compound is *.

• 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

SECTION10.4

Empirical and Molecular Formulas

Molecular Formula

• The molecular formula specifies *.

• Molecular formula is always a whole-number multiple of the empirical formula.

SECTION10.4

Empirical and Molecular Formulas

Molecular Formula (cont.)

SECTION10.4

Empirical and Molecular Formulas

• 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.

SECTION10.5

Formulas of Hydrates

Naming Hydrates• A hydrate is a *

• The number of water molecules associated with each formula unit of the compound is written following a dot.

• Sodium carbonate decahydrate = Na2CO3 • 10H2O

SECTION10.5

Formulas of Hydrates

Naming Hydrates (cont.)

SECTION10.5

Formulas of Hydrates

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.

SECTION10.5

Formulas of Hydrates

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.

SECTION10.5

Formulas of Hydrates

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.

SECTION10.5

Formulas of Hydrates

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.

SECTION10.1

Measuring Matter

Study Guide

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.

SECTION10.2

Mass and the Mole

Study Guide

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.

SECTION10.3

Moles of Compounds

Study Guide

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.

SECTION10.4

Empirical and Molecular Formulas

Study Guide