Chapter 3. Molecules, Compounds and Chemical Equations

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Chapter 3. Molecules, Compounds and

Chemical Equations

Modified by Dr. Cheng-Yu Lai

Stoichiometry

• Mole concept and Avogadro’s Number • Determining Chemical Formulas

• Name Compound

• Balancing Chemical Reactions • Yields

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An Atomic-Level View of Elements and Compounds

• Elements may be either atomic or molecular.

• Compounds may be either molecular or ionic.

single atoms

Chemical bonds between atoms

molecular or ionic compound.

multiple atoms

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View of Elements

• Atomic elements

exist in nature with

single

atoms

as their basic units. Most elements

fall into this category.

• Examples are Na, Ne, C, K, Mg, etc.

• Molecular elements

do not normally exist

in nature with single atoms as their basic

units; instead, they exist as

molecules—two

or more atoms of the element bonded

together.

• There only seven diatomic elements and they are H₂, N₂, O₂, F₂, Cl₂, Br₂, and I₂.

• Also, P₄ and S₈ are polyatomic elements.

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• Elements may be either atomic or molecular. • Compounds may be either molecular or ionic.

single atoms

Chemical bonds between atoms

molecular or ionic compound.

multiple atoms

An Atomic-Level View of Elements and Compounds

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Molecules and Chemical Bonds

Covalent Bond: A bond that results when two atoms share several (usually two) electrons. Typically a nonmetal

bonded to a nonmetal.

Molecule: The unit of matter that results when two or more atoms are joined by covalent bonds.

Based on atom- atom

interactions-Chemical bonds are classified into two types:

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Molecules Compounds, Covalent Bonds

To visualize the molecules, it helps to imagine the

individual atoms as spheres jointed together to form molecules with specific 3-D shapes shown below.

A ball-and-stick molecular model represents atoms as balls and chemical bonds as sticks; how the two connect reflects a molecule’s shape.

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Which one is element ? Molecular element ? compound

Chapter 3/7 2CO₂(l)

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Ionic Compound and Ionic Bonds

• Ionic Bond: Electrostatic attraction between the

resulting charged particles after a complete transfer of one or more electrons from one atom to another;

typically a metal bonded to a nonmetal. Ion: A charged particle, could be :

1. Cation: A positively charged particle. Metals tend to lose electron(s) and form cations.

2. Anion: A negatively charged particle. Nonmetals tend to gain electron(s) and form anions.

11 protons 11 electrons 17 protons 17 electrons Na+ ₊ _Cl– Na + Cl₂ 11 protons 10 electrons 17 protons 18 electrons Chapter 3/8

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Ions and Ionic Bonds

Na+ _{+ Cl}

-Na + Cl₂

In the formation of sodium chloride, one electron is transferred from the sodium atom to the chlorine atom.

2 1

 Cation: A positive ion -Na+

 Anion: A negative ion -Cl−,

 Ionic Bonding: Force of

attraction between oppositely charged ions.

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Predicting Ionic Charges

Group 1: Lose 1 electron to form 1+ ions

H+ _Li+ _Na+ _K+ _Rb+ _Cs+

Chapter 3/10

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Predicting Ionic Charges

Group 2: Loses 2 electrons to form 2+ ions

Be2+ _Mg2+ _Ca2+ _Sr2+ Ba2+

Chapter 3/11

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Predicting Ionic Charges

Group 13: Loses 3 electrons to form 3+ ions

B

3+

Al

3+

Ga

3+ Chapter 3/12

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Predicting Ionic Charges

Group 14: Loses 4 electrons or gains 4 electrons Chapter 3/13

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Predicting Ionic Charges

Group 15: Gains 3 electrons to form 3- ions N 3-P 3-As 3-Nitride Phosphide Arsenide Chapter 3/14

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Predicting Ionic Charges

Group 16: Gains 2 electrons to form 2- ions O 2-S 2-Se 2-Oxide Sulfide Selenide Chapter 3/15

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Predicting Ionic Charges

Group 17: Gains 1 electron to form 1- ions F 1-Cl 1-Br 1-Fluoride Chloride Bromide I1- _Iodide Chapter 3/16

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Representing Compounds with its

Chemical Formulas

• Chemical formula indicates the elements

present in the compound and the relative

number of atoms or ions of each.

– Water is represented as H₂O.

– Carbon dioxide is represented as CO₂.

– Sodium Chloride is represented as NaCl.

– Carbon tetrachloride is represented as CCl₄.

Chapter 3/18

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Types of Chemical Formulas

• Chemical formulas can generally be

categorized into three different types:

• Empirical formula

• Molecular formula

• Structural formula

Chapter 3/19

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Types of Chemical Formulas

• An empirical formula gives the relative

number of atoms of each element in a

compound.

• A molecular formula gives the actual

number of atoms of each element in a

molecule of a compound.

(a) For C₄H₈, the greatest common factor is 4. The empirical formula is therefore CH₂.

(b) For B₂H₆, the greatest common factor is 2. The empirical formula is therefore BH₃.

(c) For CCl₄, the only common factor is 1, so the empirical formula and the molecular formula

are identical. _Chapter

(21)

Types of Chemical Formulas

• A structural formula uses lines to

represent covalent bonds and shows how

atoms in a molecule are connected or

bonded to each other. The structural

formula for H

₂

O

₂

is shown below ( seldom

used in 100 level class )

Chapter 3/21

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Formulas

Formulas for molecular compounds MIGHT

be empirical (lowest whole number ratio).

Molecular: H₂O C₆H₁₂O₆ C₁₂H₂₂O₁₁ Empirical: H₂O CH₂O C₁₂H₂₂O₁₁ Empirical Formula

Simplest, whole-number ratio of the atoms of elements in a compound

Chapter 3/22

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Chemistry: A Molecular Approach, 3rd Edition

Nivaldo J. Tro

Formula Mass

•

The mass of an individual molecule or formula unit

 also known as molecular mass or molecular weight

•

Sum of the masses of the atoms in a single molecule or formula unit

whole = sum of the parts!

• Mass of 1 molecule of H

₂

O

= 2(1.01 g/mole H) + 16.00g/mole O = 18.02 g/mole

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Calculating Formula Mass

Calculate the formula mass of 1 mole of magnesium carbonate, MgCO₃.

24.31 g + 12.01 g + 3(16.00 g) = 84.32 g/mole Its the sum of the individual atomic masses of each atom constituting the molecule.

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Calculating Percentage Composition

Calculate the percentage composition of magnesium carbonate, MgCO₃.

From previous slide:

24.31 g + 12.01 g + 3(16.00 g) = 84.32 g 100.00 24.31 100 28.83% 84.32 Mg  _ _    

12.01 100 14.24%

84.32 C



















48.00

100 56.93%

84.32 O



















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Magnesium Carbonate contains 28.83% Mg, 14.24% C, and 56.93% O by mass. What is the empirical formula of magnesium carbonate?

Empirical Formula Determination

Empirical Formula

Simplest, whole-number ratio of the atoms of elements in a compound Can be determined from elemental analysis

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Empirical Formula Determination

1. Base calculation on 100 grams of

compound.

2. Determine moles of each element in

100 grams of compound.

3. Divide each value of moles by the

smallest of the values.

4. Multiply each number by an integer

to obtain all whole numbers.

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Magnesium Carbonate contains 28.83% Mg, 14.24% C, and 56.93% O by mass. What is the empirical formula of magnesium carbonate?

Empirical Formula Determination

185 . 1 31 . 24 83 . 28 _  g Mg 185 . 1 01 . 12 24 . 14   g C 558 . 3 16 93 . 56   g O

1. Base calculation on 100 grams of compound.

2. Determine moles of each element in 100 grams of compound.

3. Divide each value of moles by the smallest of the values. 4. Multiply each number by an

integer to obtain all whole numbers.

Mg

_1.185

C

_1.185

O

_3.558 3 1 1 185 . 1 558 . 3 185 . 1 185 . 1 185 . 1 185 . 1

O

C

Mg

O

C

Mg



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Empirical Formula Determination

Adipic acid contains 49.32% C, 43.84% O, and 6.85% H by mass. What is the empirical formula of adipic acid?











49.32g C 1 mol C =4.107 mol C 12.01 g C











6.85 1 6.78 1.01 g H mol H mol H g H 











43.84 1 2.74 16.00 g O mol O mol O g O  Chapter 3/29

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Empirical Formula Determination

(part 2)

Divide each value of moles by the smallest of the values. Carbon: Hydrogen: Oxygen:

4.107

1.50

2.74 mol C

mol O



6.78

2.47

2.74 mol H

mol O



2.74

1.00

2.74 mol O

mol O



C

_1.5

H

_2.47

O

₁ Chapter 3/30

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Empirical Formula Determination

(part 3)

Multiply each number by an integer to obtain all whole numbers.

Carbon: 1.50 Hydrogen: 2.50 Oxygen: 1.00

x 2 x 2 x 2

3 5 2

Empirical formula:

C

₃

H

₅

O

₂

Empirical Formula

Simplest, whole-number ratio of the atoms of elements in a compound Can be determined from elemental analysis

Chapter 3/31

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Finding the Molecular Formula

The empirical formula for adipic acid is

C₃H₅O₂. The molecular mass of adipic acid is 146 g/mol. What is the molecular

formula of adipic acid?

1. Find the formula mass of C₃H₅O₂

3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g

Chapter 3/32

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Finding the Molecular Formula

3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g

2. Divide the molecular mass by the mass given by the emipirical formula.

146

2

73 

The molecular formula is a multiple

of the empirical formula

.

_Chapter

(34)

Finding the Molecular Formula

3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g

3. Multiply the empirical formula by this number to get the molecular formula.

(C

₃

H

₅

O

₂

) x 2 =

C

₆

H

₁₀

O

₄

146

2

73 

Molecular formula = (empirical formula)n,

where n is a positive integer.

Chapter 3/34

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Empirical Formula Determination

Adipic acid contains 26.0% N and 74.0% O by mass. What is the empirical formula of a compound?

(26.0 g N )( 1 mole N) --- = 1.857 mole N 14.00 g N (74.0 g O )( 1 mole O) --- = 4.628 mole O 15.99 g O

Divide each value of moles by the smallest of the values.

N

_1.O

O

_2.5

=

N

₂

O

₅

Multiply each number by an integer to

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Moles and Formula Mass

mole compound

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Instructor Resource DVD for Chemistry, 6th Edition

John McMurry & Robert C. Fay

Worked Example 3.6 Converting Moles To Mass

Strategy

The problem gives the number of moles of NaHCO₃and asks for a mole-to-mass conversion. First, calculate the molar mass of NaHCO₃. Then use molar mass as a conversion factor, and set up an equation so that the unwanted unit cancels.

Solution

Formula mass of NaHCO₃ = 23.0 amu + 1.0 amu + 12.0 amu + (3 × 16.0 amu) = 84.0 amu

Molar mass of NaHCO₃= 84.0 g/mol

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Nivaldo J. Tro

Example 3.13 The Mole Concept—Converting between Mass and Number of Molecules

An aspirin tablet contains 325 mg of acetylsalicylic acid (C₉H₈O₄). How many acetylsalicylic acid molecules does it contain?

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Nivaldo J. Tro

Example 3.14 Mass Percent Composition

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Nivaldo J. Tro

Example 3.18 Obtaining an Empirical Formula from Experimental Data

A laboratory analysis of aspirin determined the following mass percent composition: C 60.00%

H 4.48% O 35.52% Find the empirical formula.

Given: In a 100 g sample: 60.00 g C, 4.48 g H,

35.52 g O

Find: empirical formula

Step 3 Write down a pseudoformula for the compound using the number of moles of each element (from step 2) as subscripts.

C_4.996H_4.44O_2.220

The correct empirical formula is C₉H₈O₄.

What is empirical formula of a compound that is 30.4 % of nitrogen and 69.6 % of oxygen ? Exam !

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Nivaldo J. Tro

Example 3.19 Calculating a Molecular Formula from an Empirical Formula and Molar Mass

Sort

Butanedione—a main component responsible for the smell and taste of butter and cheese—contains the elements carbon, hydrogen, and oxygen. The empirical formula of butanedione is C₂H₃O, and its molar mass is 86.09 g/mol. Find its molecular formula.

You are given the empirical formula and molar mass of butanedione and asked to find the molecular formula.

Given: Empirical formula = C₂H₃O molar mass = 86.09 g/mol

Find: molecular formula

Strategize Solve

Calculate the empirical formula mass.

Divide the molar mass by the empirical formula mass to find n.

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Chemical Equations

• Provide information about the reaction

– Formulas of reactants and products – States of reactants and products

– Relative numbers of reactant and product molecules that are required

– Can be used to determine weights of reactants used and products that can be made

Chapter 3/42

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How to Balance Chemical Equations

A balanced chemical equation shows that the law of

conservation of mass is adhered to.

In a balanced chemical equation, the numbers and kinds of atoms on both sides of the reaction arrow are identical. 2NaCl(s) 2Na(s) + Cl₂(g) right side: 2 Na 2 Cl left side: 2 Na 2 Cl

Chemical Equation - Shorthand way of describing a reaction

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Count Atoms

• 3H₂O

• subscripts – little numbers that tell how many atoms there are (ex: In 3H₂O, the 2 is the subscript)

• coefficients – regular-sized numbers that tell how many molecules there are (ex: In 3H₂O, the 3 is the coefficient)

(45)

Nivaldo J. Tro

Example 3.22 Balancing Chemical Equations

How many carbon atoms are found in 10 g of C₂H₅OH ? Molar Mass of C₂H₅OH=46 , Exam = 10.0 g C2H5OH 46.0 g C2H5OH 1 mol C2H5OH 1 mol C2H50H 2 mole Carbon 1 mol 6.02X1023 _atoms 2.62X1023 atoms

Balancing Equations Worksheet

____ C3H7OH + ____ O2  ____ CO2 + ____ H2O ____ N₂ + ____ H₂  ____ NH₃

1 N₂ + 3 H₂  2 NH₃

____ CH₄ + ____ O₂  ____ CO₂ + ____ H₂O 1 CH₄ + 2 O₂  1 CO₂ + 2 H₂O

(46)

Ionic Compound Formulas

Formulas for ionic compounds are ALWAYS

empirical (lowest whole number ratio). Examples:

NaCl MgCl₂ Al₂(SO₄)₃ K₂CO₃

Chapter 3/46

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Cations

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Anions

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Writing

Ionic Compound Formulas

Example: Barium nitrate

1. Write the formulas for the cation and anion, including CHARGES!

Ba

2+

NO

3

-2. Check to see if charges are balanced.

3. Balance charges , if necessary, using subscripts.

Use parentheses if you need more than one of a

polyatomic ion. Not balanced

(

)

₂

Chapter 3/49

(50)

Writing

Ionic Compound Formulas

Example: Iron(III) chloride

Fe

3+

Cl

3

Chapter 3/50

(51)

Writing

Ionic Compound Formulas

Example: Magnesium carbonate

Mg

2+

_CO

3

2-2. Check to see if charges are balanced.

They are balanced

Chapter 3/51

(52)

Writing

Ionic Compound Formulas

Example: Zinc hydroxide

Zn

2+

OH

(

)

₂

Chapter 3/52

(53)

Writing

Ionic Compound Formulas

Example: Aluminum phosphate

Al

3+

_PO

4

3-2. Check to see if charges are balanced.

They ARE balanced

Chapter 3/53

(54)

Naming Binary Molecular Compounds

 Compounds between two nonmetals

 First element in the formula is named first.  Keeps its element name

 Gets a prefix if there is a subscript on it  Second element is named second

 Use the root of the element name plus the

-ide suffix

 Always use a prefix on the second element

Chapter 3/54

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Naming

Ionic Compounds

Naming

Ionic Compounds



Cation first, then anion



Monatomic cation = name of the

element



Ca

2+

= calcium ion



Monatomic anion =

root

+

-ide



Cl

-

=

chloride



CaCl

₂

= calcium chloride



Cation first, then anion



Monatomic cation = name of the

element



Ca

2+

= calcium ion



Monatomic anion =

root

+

-ide



Cl

-

=

chloride



CaCl

₂

= calcium chloride

Chapter 3/55

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Naming Ionic Compounds

 some metal forms more than one cation

 use Roman numeral in (in parentheses) that

indicates the charge of the metal in that particular compound.

 PbCl₂

 Pb2+ is cation

 PbCl₂ = lead(II) chloride

 some metal forms more than one cation

 use Roman numeral in (in parentheses) that

indicates the charge of the metal in that particular compound.

 PbCl₂

 Pb2+ is cation

 PbCl₂ = lead(II) chloride

Metals with multiple oxidation states

For example, we distinguish between Fe2+ and Fe3+ as follows:

Fe2+ Iron(II)

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Naming Chemical Compounds

Because nonmetals often combine with one another in different proportions to form

different compounds, numerical prefixes are usually included in the names of binary molecular compounds.

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Naming Binary Compounds

P

₂

O

₅

=

CO

₂

=

N

₂

O =

di

phosphorus

pent

oxide

carbon

di

oxide

di

nitrogen

mon

oxide

SO

₃

= Sulfur

tri

oxide

Lead(II)

mon

oxide

PbO =

Lead(II)

di

oxide

PbO

₂

=

Chapter

(59)

Practice – Write the Formula

Compound Name Compound Formula

Carbon dioxide Carbon monoxide Diphosphorus pentoxide Sulfur trioxide Copper(II) oxide Carbon tetrabromide Lead (II) Oxide

Lead (IV) dioxide Iodine trichloride Sodium nitride

copper(II) phosphate

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Answers – Write the Formula

Compound Name Compound Formula

Carbon dioxide CO₂

Carbon monoxide CO

Diphosphorus pentoxide P₂O₅

Sulfur trioxide SO₃

Copper(II) oxide CuO

Carbon tetrabromide CBr₄

Lead (II) oxide PbO

Lead (IV) dioxide PbO₂ Iodine trichloride ICl₃

Sodium nitride NaN₃

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Hydrated Ionic Compounds

• Hydrates are ionic compounds containing

a specific number of water molecules

associated with each formula unit.

– For example, the formula for epsom salts is

MgSO₄ • 7H₂O.

– Its systematic name is magnesium sulfate

heptahydrate.

– CoCl₂ • 6H₂O is cobalt(II)chloride hexahydrate.

Chapter 3/61

(62)

Other common hydrated ionic compounds and their

names are as follows:

– CaSO₄ • 1/2H₂O is called

calcium sulfate hemihydrate. – BaCl₂ • 6H₂O is called barium

chloride hexahydrate.

– CuSO₄ • 6H₂O is called copper

sulfate hexahydrate.

Hydrates

Chapter 3/62

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Acids

• Acids are molecular compounds that

release hydrogen ions (H

+

) when dissolved

in water.

• Acids are composed of hydrogen, usually

written first in their formula, and one or

more nonmetals, written second.

– HCl is a molecular compound that, when

dissolved in water, forms H+_(aq) and Cl–_(aq) ions,

where aqueous (aq) means dissolved in water.

Chapter 3/63

(64)

•

Binary acids have H+1

cation and nonmetal anion.

Acids

Chapter 3/64

•

Oxyacids have H+ cation and oxyanion ( an

(65)

Naming Binary Acids

• Write a hydro- prefix.

• Follow with the nonmetal name.

• Change ending on nonmetal name to –ic.

• Write the word acid at the end of the name.

Chapter 3/65 HF – hydrofluoric acid HCl – hydrochloric acid HBr – hydrobrmoic acid HI – hydroiodic acid

(66)

Naming Oxyacids

• If polyatomic ion name ends in –ate, then change

ending to –ic suffix.

HNO₃ (oxyanion is nitrate , NO₃ - ) = Nitric acid

H₂SO₄ (oxyanion is sulfate , SO₄ 2- ) = Sulfuric acid

• If polyatomic ion name ends in –ite, then change

ending to –ous suffix.

HNO₂ (oxyanion nitrite , NO₂ - ) = Nitrous acid

H₂SO₃ (oxyanion is sulfite , SO₃ 2- _{) = Sulfurous acid}

(67)

Carbonic acid

Hydrochloric acid

acetic acid

1. H

₂

CO3- Oxyacids

2. HCl -Binary acids

3. HC

₂

H

₃

O

₂

- Oxyacids

Name the Following Acids

Chapter 3/67

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Nivaldo J. Tro

Acid Rain and Climate Change

• Certain pollutants—such as NO, NO₂, SO₂, SO₃—form acids

when mixed with water, resulting in acidic rainwater.

• Acid rain can fall or flow into lakes and streams, making these

bodies of water more acidic.

• The details on this marble statue have been eaten over the years by

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Nivaldo J. Tro

Organic Compounds

• Early chemists divided compounds into two

types: organic and inorganic.

• Compounds from living things were called

organic; compounds from the nonliving

environment were called inorganic.

• Organic compounds are easily decomposed and

could not be made in the lab.

• Inorganic compounds are very difficult to

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Nivaldo J. Tro