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Chapter 2

Atoms, Molecules,

and Ions

Section 2.1

The Early History of Chemistry

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(2.1) The early history of chemistry
(2.2) Fundamental chemical laws
(2.3) Dalton’s atomic theory
(2.4) Early experiments to characterize the atom
(2.5) The modern view of atomic structure: An introduction
(2.6) Molecules and ions
(2.7) An introduction to the periodic table
(2.8) Naming simple compounds

Chapter 2

Table of Contents

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Early History of Chemistry

Applications of chemistry before 1000 B.C

Usage of embalming fluids
Production of metals for weapons and ornaments

The Greeks (400 B.C)

Proposed that matter was composed of earth, fire, air, and water
Questioned whether matter is infinitely divisible or is composed of small, indivisible particles

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Section 2.1

The Early History of Chemistry

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Early History of Chemistry - Alchemy

Alchemists dominated the field of chemistry for 2000 years

Helped discover several elements
Learned to prepare mineral acids

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Section 2.1

The Early History of Chemistry

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Modern Chemistry

Foundation was laid by:

Georg Bauer, who developed systematic metallurgy
Paracelsus, who discovered the medicinal applications of minerals

Robert Boyle

Performed quantitative experiments to measure the relationship between the pressure and volume of air

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Section 2.1

The Early History of Chemistry

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Modern Chemistry (Continued 1)

Developed the first experimental definition of an element

A substance is an element unless it can be broken down into two or more simpler substances

Held on to certain alchemists’ views

Metals are not true elements
Eventually, a method to change one metal to another will be found

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Section 2.1

The Early History of Chemistry

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Modern Chemistry (Continued 2)

17^th and 18^th century

Rise in interest in the phenomenon of combustion
Georg Stahl suggested that a substance called phlogiston flowed out of burning material

Substances that burn in a closed container eventually stop burning since the air in the container is saturated with phlogiston

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Section 2.1

The Early History of Chemistry

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Modern Chemistry (Continued 3)

Joseph Priestley discovered that oxygen vigorously supported combustion

Oxygen was supposed to be low in phlogiston

Was originally called dephlogisticated air

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Section 2.1

The Early History of Chemistry

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Antoine Lavoisier

Proposed the law of conservation of mass

Law of conservation of mass: Mass is neither created nor destroyed in a chemical reaction

Showed that combustion involves oxygen, not phlogiston
Discovered that life is supported by a process that involves oxygen and is similar to combustion

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Section 2.2

Fundamental Chemical Laws

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Joseph Proust

Proposed the Proust’s law or law of definite proportion

Law of definite proportion: A given compound always contains exactly the same proportion of elements by mass

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Section 2.2

Fundamental Chemical Laws

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John Dalton

Suggested that elements were composed of tiny individual particles

A given compound always contains the same combination of these atoms

Proposed the law of multiple proportions

Law of multiple proportions: When two elements form a series of compounds:

The ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers

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Section 2.2

Fundamental Chemical Laws

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Example 2.1 - Illustrating the Law of Multiple Proportions

The following data were collected for several compounds of nitrogen and oxygen:

Show how these data illustrate the law of multiple proportions

Section 2.2

Fundamental Chemical Laws

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Example 2.1 - Solution

For the law of multiple proportions to hold, the ratios of the masses of nitrogen combining with 1 g of oxygen in each pair of compounds should be small whole numbers

Therefore, compute the ratios as follows:

Section 2.2

Fundamental Chemical Laws

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Example 2.1 - Solution (Continued)

These results support the law of multiple proportions

Section 2.2

Fundamental Chemical Laws

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Dalton’s Atomic Theory

Each element is made up of tiny particles called atoms
Atoms of a given element are identical

Atoms of different elements are different in some fundamental way or ways

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Section 2.3

Dalton’s Atomic Theory

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Dalton’s Atomic Theory (Continued)

Chemical compounds are formed when atoms of different elements combine with each other

A given compound always has the same relative numbers and types of atoms

Chemical reactions involve reorganization of the atoms

The atoms themselves are not changed in a chemical reaction

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Section 2.3

Dalton’s Atomic Theory

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Table of Atomic Masses

Dalton prepared the first table of atomic masses (atomic weights)
Assumption - Nature is as simple as possible

Many masses were proved to be wrong due to the assumption

Section 2.3

Dalton’s Atomic Theory

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Gay-Lussac

Measured the volumes of gases that reacted with each other under the same temperature and pressure

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Section 2.3

Dalton’s Atomic Theory

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Avogadro’s Hypothesis

At the same temperature and pressure, equal volumes of different gases contain the same number of particles

Volume of a gas is determined by the number, not the size, of molecules

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The spheres represent atoms in the molecules

Section 2.3

Dalton’s Atomic Theory

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Combining Gay-Lussac’s Result and Avogadro’s Hypothesis

If Avogadro’s hypothesis is correct,

Can be expressed as:

Section 2.3

Dalton’s Atomic Theory

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J. J. Thomson

Studied electric discharge using cathode-ray tubes

Cathode-ray tubes: Partially evacuated tubes
When high voltage was applied to the tube, a cathode ray was produced

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Section 2.4

Early Experiments to Characterize the Atom

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Figure 2.8 - Deflection of Cathode Rays by an Applied Electrical Field

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The ray was produced at the negative electrode and was repelled by the negative pole of an applied electric field

Section 2.4

Early Experiments to Characterize the Atom

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J. J. Thomson - Contributions

Postulated the existence of negatively charged particles (electrons)
Determined the charge-to-mass ratio of an electron

e - Charge on the electron (in coulombs)
m - Electron mass (in grams)

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Section 2.4

Early Experiments to Characterize the Atom

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J. J. Thomson - The Structure of Atoms

Thomson attempted to understand the structure of an atom
Assumptions

All atoms contain electrons as electrons can be produced from electrodes made of various metals
Atoms must contain some amount of positive charge

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Section 2.4

Early Experiments to Characterize the Atom

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J. J. Thomson’s Plum Pudding Model

Atoms consist of a diffuse cloud of positive charge

Negative electrons are randomly embedded in it

Section 2.4

Early Experiments to Characterize the Atom

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Robert Millikan

Performed experiments involving charged oil drops
Determined the magnitude of the charge on a single electron
Calculated the mass of an electron

9.11 ×10^–31 kg

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Section 2.4

Early Experiments to Characterize the Atom

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Radioactivity

Henri Becquerel

Discovered radioactivity by observing the spontaneous emission of radiation by uranium

Types of radioactive emission

Gamma rays (γ) - High-energy light
Beta particles (β) - High-speed electrons
Alpha particles (α) - Particles with a 2+ charge

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Section 2.4

Early Experiments to Characterize the Atom

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Rutherford’s Experiment

Carried out to test the accuracy of Thomson’s plum pudding model
Involved directing α particles at a thin sheet of metal foil

Expectation - α particles will pass through the foil with minor deflections in path

Section 2.4

Early Experiments to Characterize the Atom

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Rutherford’s Experiment - Results

Most α particles passed through the foil
Many particles were deflected at large angles
Some particles were reflected

Particles did not hit the detector

Section 2.4

Early Experiments to Characterize the Atom

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Figure 2.12 - Rutherford’s Experiment on α-Particle Bombardment of a Metal Foil

Section 2.4

Early Experiments to Characterize the Atom

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Rutherford’s Experiment - Conclusions

Large deflections of the α particles are caused by a center of concentrated positive charge

The center contains most of the atom’s mass
Deflected α particles had a close encounter with the massive positive center of the atom

Reflected α particles made a direct hit on the massive positive center

Most α particles pass directly through the foil because the atom is mostly open space

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Section 2.4

Early Experiments to Characterize the Atom

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Rutherford’s Experiment - Conclusions (Continued)

Nuclear atom has a dense center of positive charge (nucleus)
Electrons travel around the nucleus at a large distance relative to the nucleus

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Section 2.4

Early Experiments to Characterize the Atom

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Figure 2.13 - Rutherford’s Experiment

The expected results of the metal foil experiment if Thomson’s model were correct

Actual results

Section 2.4

Early Experiments to Characterize the Atom

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Critical Thinking

You have learned about three different models of the atom - Dalton’s model, Thomson’s model, and Rutherford’s model

What if Dalton was correct? What would Rutherford have expected from his experiments with gold foil?
What if Thomson was correct? What would Rutherford have expected from his experiments with gold foil?

Section 2.4

Early Experiments to Characterize the Atom

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Atomic Structure

Electrons are negatively charged particles that are found outside the nucleus
The nucleus contains:

Protons: Contain positive charge that is equal in magnitude to the electron’s negative charge
Neutrons: Contain no charge and have virtually the same mass as a proton

Atoms of different elements show different chemical behavior

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Section 2.5

The Modern View of Atomic Structure: An Introduction

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Table 2.1 - The Mass and Charge of the Electron, Proton, and Neutron

Section 2.5

The Modern View of Atomic Structure: An Introduction

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The Nucleus

Small compared to the overall size of the atom
High in density

Accounts for almost all of the atom’s mass

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Section 2.5

The Modern View of Atomic Structure: An Introduction

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Isotopes

Atoms with the same number of protons but different numbers of neutrons
Depict almost identical chemical properties
In nature, most elements contain mixtures of isotopes

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Section 2.5

The Modern View of Atomic Structure: An Introduction

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Figure 2.15 - Two Isotopes of Sodium

Section 2.5

The Modern View of Atomic Structure: An Introduction

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Identifying Isotopes

Atomic number (Z): Number of protons

Written as a subscript

Mass number (A): Total number of protons and neutrons

Written as a superscript

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Mass number

Atomic number

Element symbol

Section 2.5

The Modern View of Atomic Structure: An Introduction

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Critical Thinking

The average diameter of an atom is 2×10^–10 m

What if the average diameter of an atom were 1 cm?
How tall would you be?

Section 2.5

The Modern View of Atomic Structure: An Introduction

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Interactive Example 2.2 - Writing the Symbols for Atoms

Write the symbol for the atom that has an atomic number of 9 and a mass number of 19

How many electrons and how many neutrons does this atom have?

Section 2.5

The Modern View of Atomic Structure: An Introduction

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Interactive Example 2.2 - Solution

The atomic number 9 means the atom has 9 protons

This element is called fluorine, symbolized by F
The atom is represented as follows:

The atom is called fluorine nineteen

Section 2.5

The Modern View of Atomic Structure: An Introduction

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Interactive Example 2.2 - Solution (Continued)

Since the atom has 9 protons, it also must have 9 electrons to achieve electrical neutrality
The mass number gives the total number of protons and neutrons, which means that this atom has 10 neutrons

Section 2.5

The Modern View of Atomic Structure: An Introduction

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Exercise

How many protons and neutrons are in the nucleus of each of the following atoms?

In a neutral atom of each element, how many electrons are present?

⁷⁹Br
⁸¹Br
²³⁹Pu
¹³³Cs

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35 p, 44 n, 35 e

35 p, 46 n, 35 e

94 p, 145 n, 94 e

55 p, 78 n, 55 e

Section 2.5

The Modern View of Atomic Structure: An Introduction

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Chemical Bonds

Forces that hold atoms together in a compound
Can be formed by sharing of electrons

Leads to the formation of a covalent bond
Molecule: Collection of atoms

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Section 2.6

Molecules and Ions

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Chemical Formula

Symbols of elements used to indicate types of atoms present in the molecule

Subscript indicates the relative number of atoms
Example - Formula for carbon dioxide is CO₂

Implies that each molecules of CO₂ contains one atom of carbon and two atoms of oxygen

Section 2.6

Molecules and Ions

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Methods of Representing Molecules

Structural formula: Depicts individual bonds in a molecule

May or may not indicate the actual shape of the molecule

Space-filling model: Illustrates the relative sizes of atoms and their relative orientations in a molecule
Ball-and-stick model

Section 2.6

Molecules and Ions

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Figure 2.16 - Structure of Methane

Structural formula

Space-filling model

Ball-and-stick model

Section 2.6

Molecules and Ions

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Ion

Atom or group of atoms with a net positive or negative charge

Chemical bonds can be a result of ionic attraction
Cation: Positive ion formed by losing electrons
Anion: Negative ion formed by gaining electrons

Ionic bonding: Force of attraction between oppositely charged ions

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Section 2.6

Molecules and Ions

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Ionic Solids

Solids containing oppositely charged ions
Can consist of:

Simple ions

Example - Sodium chloride (common salt)

Polyatomic ions: Composed of many atoms

Example - Ammonium nitrate (NH₄NO₃)

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Section 2.6

Molecules and Ions

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Exercise

Would you expect each of the following atoms to gain or lose electrons when forming ions?

What ion is the most likely in each case?

Ra
In
P
Te

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Loses 2 e^–to form Ra²⁺

Loses 3 e^–to form In³⁺

Gains 3 e^–to form P^3–

Gains 2 e^–to form Te^2–

Section 2.6

Molecules and Ions

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The Periodic Table

A chart that provides information about elements

Letters in boxes are symbols of elements
Number above every symbol is the element’s atomic number

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Atomic number

Element symbol (Hydrogen)

Section 2.7

An Introduction to the Periodic Table

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Structure of the Periodic Table

Groups or families: Elements in the vertical columns with similar chemical properties

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Section 2.7

An Introduction to the Periodic Table

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Structure of the Periodic Table (Continued)

Periods: Horizontal rows of elements

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Section 2.7

An Introduction to the Periodic Table

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Table of Common Charges

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Group or family	Charge
Alkali metals (1A)	1+
Alkaline earth metals (2A)	2+
Halogens (7A)	1–
Noble gases (8A)	0

Section 2.7

An Introduction to the Periodic Table

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Metals and Nonmetals

Metals

Lose electrons to form positive ions
Efficient conductors of heat and electricity
Malleable, ductile, and have lustrous appearance

Nonmetals

Gain electrons to form negative ions
Appear in the upper-right corner of the periodic table
Form covalent bonds

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Section 2.7

An Introduction to the Periodic Table

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Binary Compounds

Composed of two elements
Include covalent and ionic compounds

Binary ionic compounds: Contain a cation, which is written first in the formula, and an anion

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Section 2.8

Naming Simple Compounds

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Naming Binary Ionic Compounds (Type I)

The cation is always named first and the anion second
A monatomic cation takes its name from the name of the parent element
A monatomic anion is named by taking the root of the element name and adding –ide

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Section 2.8

Naming Simple Compounds

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Table 2.3 - Common Monatomic Cations and Anions

Section 2.8

Naming Simple Compounds

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Interactive Example 2.3 - Naming Type I Binary Compounds

Name each binary compound

CsF
AlCl₃
LiH

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Section 2.8

Naming Simple Compounds

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Interactive Example 2.3 - Solution

CsF is cesium fluoride
AlCl₃ is aluminum chloride
LiH is lithium hydride
Notice that, in each case, the cation is named first and then the anion is named

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Section 2.8

Naming Simple Compounds

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Binary Ionic Compounds (Type II)

Metals in such compounds form more than one type of positive ion
Naming

Charge on the metal ion must be specified

Roman numeral indicates the charge of the metal cation

Compounds containing transition metals usually require a Roman numeral
Elements that form only one cation do not need to be identified by a Roman numeral

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Section 2.8

Naming Simple Compounds

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Common Metals That Do Not Require a Roman Numeral

Group 1A elements
Group 2A elements
Aluminum
Silver
Zinc

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Section 2.8

Naming Simple Compounds

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Critical Thinking

We can use the periodic table to tell us something about the stable ions formed by many atoms

For example, the atoms in column 1 always form 1+ ions
The transition metals, however, can form more than one type of stable ion
What if each transition metal ion had only one possible charge?
How would the naming of compounds be different?

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Section 2.8

Naming Simple Compounds

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Figure 2.20 - Flowchart for Naming Binary Ionic Compounds

Section 2.8

Naming Simple Compounds

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Interactive Example 2.6 - Naming Binary Compounds

Give the systematic name for each of the following compounds:

CoBr₂
CaCl₂

Given the following systematic names, write the formula for each compound:

Chromium(III) chloride
Gallium iodide

Section 2.8

Naming Simple Compounds

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Interactive Example 2.6 - Solution (1)

Formula	Name	Comments
CoBr₂	Cobalt(II) bromide	Cobalt is a transition metal; the compound name must have a Roman numeral The two Br^– ions must be balanced by a Co²⁺ ion
CaCl₂	Calcium chloride	Calcium, an alkaline earth metal, forms only the Ca²⁺ ion A Roman numeral is not necessary
Al₂O₃	Aluminium oxide	Aluminum forms only the Al³⁺ ion A Roman numeral is not necessary

Section 2.8

Naming Simple Compounds

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Interactive Example 2.6 - Solution (2)

Name	Formula	Comments
Chromium(III) chloride	CrCl₃	Chromium(III) indicates that Cr³⁺ is present, so 3 Cl^– ions are needed for charge balance
Gallium iodide	GaI₃	Gallium always forms 3+ ions, so 3 I^– ions are required for charge balance

Section 2.8

Naming Simple Compounds

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Polyatomic Ions

Assigned special names that must be memorized for naming compounds
Oxyanions: Anions that contain an atom of a given element and different numbers of O₂atoms

When there are two elements in the series:

Name of the element with the smaller number of O₂ atoms ends with -ite
Name of the element with the larger number of O₂atoms ends with -ate

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Section 2.8

Naming Simple Compounds

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Polyatomic Ions (Continued)

When more than two oxyanions make up a series:

Use the prefix hypo- (less than) to name members of the series with the fewest O₂ atoms
Use the prefix per- (more than) to name members of the series with the most O₂ atoms

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Section 2.8

Naming Simple Compounds

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Table 2.5 - Common Polyatomic Ions

Section 2.8

Naming Simple Compounds

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Interactive Example 2.7 - Naming Compounds Containing Polyatomic Ions

Give the systematic name for each of the following compounds:

Na₂SO₄
Mn(OH)₂

Given the following systematic names, write the formula for each compound:

Sodium hydrogen carbonate
Sodium selenate

Section 2.8

Naming Simple Compounds

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Interactive Example 2.7 - Solution (1)

Formula	Name	Comments
Na₂SO₄	Sodium sulfate
Mn(OH)₂	Manganese(II) hydroxide	Transition metal—name must contain a Roman numeral The Mn²⁺ ion balances three OH^– ions

Section 2.8

Naming Simple Compounds

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Interactive Example 2.7 - Solution (2)

Name	Formula	Comments
Sodium hydrogen carbonate	NaHCO₃	Often called sodium bicarbonate
Sodium selenate	Na₂SeO₄	Atoms in the same group, like sulfur and selenium, often form similar ions that are named similarly Thus SeO₄^2– is selenate, like SO₄^2– (sulfate)

Section 2.8

Naming Simple Compounds

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Binary Covalent Compounds (Type III)

Formed between two nonmetals
Naming binary covalent compounds

The first element in the formula is named first, using the full element name
The second element is named as if it were an anion
Prefixes are used to denote the numbers of atoms present
The prefix mono- is never used for naming the first element

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Section 2.8

Naming Simple Compounds

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Table 2.6 - Prefixes Used to Indicate Number in Chemical Names

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Section 2.8

Naming Simple Compounds

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Interactive Example 2.8 - Naming Type III Binary Compounds

Name each of the following compounds:

PCl₅
PCl₃
SO₂

From the following systematic names, write the formula for each compound:

Sulfur hexafluoride
Sulfur trioxide

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Section 2.8

Naming Simple Compounds

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Interactive Example 2.8 - Solution

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Formula	Name
PCl₅	Phosphorus pentachloride
PCl₃	Phosphorus trichloride
SO₂	Sulfur dioxide

Name	Formula
Sulfur hexafluoride	SF₆
Sulfur trioxide	SO₃

Section 2.8

Naming Simple Compounds

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Figure 2.22 - A Flowchart for Naming Binary Compounds

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Section 2.8

Naming Simple Compounds

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Acids

Molecules in which one or more H⁺ ions are attached to an anion
Naming acids

The acid is named with the prefix hydro- and the suffix -ic if the anion ends in -ide

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Section 2.8

Naming Simple Compounds

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Acids (Continued)

If the anion contains oxygen:

The suffix -ic is added to the root name if the anion name ends in -ate
The suffix -ous is added to the root name if the anion name ends in -ite

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Section 2.8

Naming Simple Compounds

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Critical Thinking

In this chapter, you have learned a systematic way to name chemical compounds

What if all compounds had only common names?
What problems would this cause?

Section 2.8

Naming Simple Compounds

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Figure 2.24 - Flowchart for Naming Acids

Section 2.8

Naming Simple Compounds

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Exercise

Name each of the following compounds:

CuI

S₄N₄

NaHCO₃

BaCrO₄

Copper(I) iodide

Tetrasulfur tetranitride

Sodium hydrogen carbonate or sodium bicarbonate

Barium chromate

Section 2.8

Naming Simple Compounds