chem16lecture_chemicalreactions

Some Types of Chemical

Reactions

Chem 16 Lecture /Nuesca/2010

2

Chapter Goals

1. The Periodic Table: Metals, Nonmetals, and Metalloids

2. Aqueous Solutions: An Introduction

3. Reactions in Aqueous Solutions

4. Oxidation Numbers

Naming Some Inorganic Compounds

5. Naming Binary Compounds

6. Naming Ternary Acids and Their Salts

Classifying Chemical Reactions

7. Oxidation-Reduction Reactions: An Introduction

8. Combination Reactions

9. Decomposition Reactions10. Displacement Reactions11. Metathesis Reactions 12. Summary of Reaction Types13. Synthesis Question

3

The Periodic Table: Metals, Nonmetals,

and Metalloids

• 1869 - Mendeleev & Meyer

– Discovered the periodic law

• The properties of the elements are periodic

functions of their atomic numbers.

4


and Metalloids

• Groups or families– Vertical group of elements on periodic table

– Similar chemical and physical properties

5


and Metalloids

• Period– Horizontal group of elements on periodic table

– Transition from metals to nonmetals

6


and Metalloids

• Some chemical properties of metals

1. Outer shells contain few electrons

2. Form cations by losing electrons

3. Form ionic compounds with nonmetals

4. Solid state characterized by metallic bonding

7


and Metalloids

• Group IA metals

– Li, Na, K, Rb, Cs, Fr

• One example of a periodic trend

– The reactions with water of Li, Na, & K

Group 1A metals are very reactive. They tend to lose electrons very quickly,

hence, in nature they exist mostly as cations (with +1 charge).

8


and Metalloids

• Group IIA metals

– alkaline earth metals

• Be, Mg, Ca, Sr, Ba, Ra

9


and Metalloids

• Some chemical properties of nonmetals

1. Outer shells contain four or more electrons

2. Form anions by gaining electrons

3. Form ionic compounds with metals and covalent

compounds with other nonmetals

4. Form covalently bonded molecules; noble gases are

monatomic

10


and Metalloids

• Group VIIA nonmetals

– halogens

– F, Cl, Br, I, At

11


and Metalloids

• Group VIA nonmetals

– O, S, Se, Te

12


and Metalloids

• Group 0 nonmetals

– noble, inert or rare gases

– He, Ne, Ar, Kr, Xe, Rn

13


and Metalloids

• Stair step function on periodic table separates

metals from nonmetals.

• Metals are to the left of

stair step.

– Approximately 80% of the

elements

• Best metals are on the far

left of the table.

14


and Metalloids



• Nonmetals are to the right of stair step. – Approximately 20% of the elements

• Best nonmetals are on the far right of the table.

15


and Metalloids



• Metalloids have one side

of the box on the stair

step.

16


and Metalloids

• Periodic trends in metallic character

PeriodicChart

More MetallicMoreMetallic

Aqueous

SolutionsReactions

18

Using Solutions in Chemical

Reactions• Titrations are a method of determining the

concentration of an unknown solutions

from the known concentration of a solution

and solution reaction stoichiometry.

– Requires special lab glassware

• Buret, pipet, and flasks

– Must have an indicator also

19


Reactions• Example 3-22: What is the molarity of a KOH

solution if 38.7 mL of the KOH solution is

required to react with 43.2 mL of 0.223 M HCl?

OH + KCl HCl + KOH 2→

20





HCl mmol 9.63 = HCl 0.223 mL 43.2

OH + KCl HCl + KOH 2

M×

→

21





KOH mmol 63.9HCl mmol 1

KOH mmol 1HCl mmol 9.63

HCl mmol 9.63 = HCl 0.223 mL 43.2


=×

×

→

M

22





KOH 249.0KOH mL 38.7

KOH mmol 9.63

KOH mmol 63.9HCl mmol 1

KOH mmol 1HCl mmol 9.63

HCl mmol 9.63 = HCl 0.223 mL 43.2


M

M

=

=×

×

→

23


Reactions• Example 3-23: What is the molarity of a barium

hydroxide solution if 44.1 mL of 0.103 M HCl is

required to react with 38.3 mL of the Ba(OH)2solution?

HCl mmol 4.54 = HCl) HCl)(0.103 mL (44.1

OH 2 + BaCl HCl 2 + Ba(OH) 222

M

→

24





HCl mmol 2

Ba(OH) mmol 1HCl mmol 54.4


OH 2 + BaCl HCl 2 + Ba(OH)

2

222

×

→

M

25





2

2

222

Ba(OH) mmol 2.27

HCl mmol 2




=

×

→

M

26





2

2

2

2

2

222

Ba(OH) 0593.0Ba(OH) mL 3.38

Ba(OH) mmol 27.2

Ba(OH) mmol 2.27

HCl mmol 2




M

M

=

=

×

→

27

Aqueous Solutions: Electrolytes.

1. Electrolytes and Extent of Ionization

• Aqueous solutions consist of a solute dissolved in water.

• Classification of solutes:

– Nonelectrolytes – solutes that do not conduct electricity in water

Examples:

C2H5OH – ethanol, C2H5OH-Methanol

Sugar solutions

Acetone, esters, many other organic compounds.

28

Aqueous Solutions: An

Introduction• C6H12O6 - glucose (blood sugar)

C

OC

C

C C

C

H

H

OH

H OH

O H

H

O H

O HH

H

H

29


Introduction• C12H22O11 - sucrose (table sugar)

C

OC

C

C C

C

H2

H

H

OH

H

O H

O H

H

O

H

OH

C C

C

O

C

CH2

H

OH

H

H

OH

C H2

OH

O H

30


Introduction• The reason nonelectrolytes do not conduct

electricity is because they do not form ions

in solution.• ions conduct electricity in solution

31

Aqueous Solutions: Strong

electrolyes: Strong Acids/Bases• Classification of solutes

– strong electrolytes - conduct electricity extremely well in dilute aqueous solutions

• Examples of strong electrolytes

1. HCl, HNO3, etc.• strong soluble acids

2. NaOH, KOH, etc.• strong soluble bases

3. NaCl, KBr, etc.• soluble ionic salts

• ionize in water essentially 100%

32

Aqueous Solutions: Weak

Electrolytes• Classification of solutes

– weak electrolytes - conduct electricity poorly

in dilute aqueous solutions

1. CH3COOH, (COOH)2• weak acids

33


electrolytes2. NH3, Fe(OH)3

• weak bases

3. some soluble covalent salts• ionize in water much less than 100%

• Weak electrolytes are poor conductors of

electricity because there are fewer ions

(charged species) present. Most of them

exist in the molecular form rather than in

ionic form.

34

Aqueous Solutions: Concept of

Acids 2. Strong and Weak Acids

• Acids are substances that generate

protons (H+ or H3O+) in aqueous solutions.

• Strong acids ionize 100% in water.

( ) ( ) ( )-

aqaq

%100

g Cl H HCl + → +≈

35

Aqueous Solutions: Acids

• Strong acids ionize 100% in water.

( ) ( ) ( )

( ) ( )-

aq3aq

OH

3

-

aq3aq3

100%

2 3

NO + H HNO

or

NO + OH OH HNO

2 +

+≈

→

→+l

No molecular HNO3 exist anymore in solution.HNO3 completely breaks up into H+ and NO3

- .

36


Introduction

Some Strong Acids and Their Anions

• Formula Name

1. HCl hydrochloric acid

2. HBr hydrobromic acid

3. HI hydroiodic acid

4. HNO3 nitric acid

5. H2SO4 sulfuric acid

6. HClO3 chloric acid

7. HClO4 perchloric acid

37


Introduction

Some Strong Acids and Their Anions

• Acid AnionName

1. HCl Cl- chloride ion

2. HBr Br- bromide ion

3. HI I- iodide ion

4. HNO3 NO3- nitrate ion

5. H2SO4 SO42- sulfate ion

6. HClO3 ClO3- chlorate ion

7. HClO4 ClO4- perchlorate ion

38

Aqueous Solutions:Weak Acids

are weak electrolytes• Weak acids ionize significantly less than

100% in water.

– Typically ionize 10% or less!

• For example acetic acid, CH3COOH, the

acid of vinegar exists primarily as a

molecule rather that as separate ions of

CH3COO- and H+. There are less ions,

hence less conductors of electricity.39


Introduction

Some Common Weak Acids and Their Anions

• Formula Name

1. HF hydrofluoric acid

2. CH3COOH acetic acid (vinegar)

3. HCN hydrocyanic acid

4. HNO2 nitrous acid

5. H2CO3 carbonic acid (soda water)

6. H2SO3 sulfurous acid

7. H3PO4 phosphoric acid

40


Introduction

Some Common Weak Acids and Their Anions

• Acid Anion Name

1. HF F- fluoride ion

2. CH3COOH CH3COO- acetate ion

3. HCN CN- cyanide ion

4. HNO2 NO2- nitrite ion

5. H2CO3 CO32- carbonate ion

6. H2SO3 SO32- sulfite ion

7. H3PO4 PO43- phosphate ion

41


Introduction3. Reversible Reactions

• As I’ve said in the previous slides,

CH3COOH acetic acid, just any weak

electrolyte, exist in solution in the

molecular form and the electrolytic form

(ions). The dissociation or ionization of

the molecule into ions is a reversible

reaction. Under normal circumstances,

it is in an equilibrium state. 42


Introduction• All weak inorganic acids ionize reversibly

or in equilibrium reactions.

– This is why they ionize less than 100%.

• CH3COOH – structure of acetic acid

C

O H

O

CH3

43


Introduction• Correct chemical symbolism for equilibrium

reactions

• Another example of a weak acid is HF acid. A

very dangerous acid which can cause serious

harm (goes into skin and dissolves the bones).

• HF (aq) < = = > H+(aq) + F

-(aq)

( ) ( )←→ +≈

aq

-

aq3

7%

3 H + COOCH COOHCH

44

Aqueous Solutions: Bases

4. Strong Bases, Insoluble Bases, and

Weak Bases (alkaline substances).

• Characteristic of common inorganic

bases is that they produce OH- ions in

solution.

• BASE (s) + H2O � OH-(aq)

45


Introduction

Common Strong Bases

• Formula Name

1. LiOH lithium hydroxide

2. NaOH sodium hydroxide

3. KOH potassium hydroxide

4. RbOH rubidium hydroxide

5. CsOH cesium hydroxide

6. Ca(OH)2 calcium hydroxide

7. Sr(OH)2 strontium hydroxide

8. Ba(OH)2 barium hydroxide

• Notice that they are all hydroxides of IA and IIA metals

46

Aqueous Solutions: Strong Bases

• Similarly to strong acids, strong bases

ionize 100% in water.

(aq)OH 2 + (aq)Ba Ba(OH)

(aq)OH + (aq)K KOH

-+2

2

-+

→

→

KOH dissolves completely in water. There are only K+ and OH- ions present in solution.

47

Aqueous Solutions: Weak Bases

Insoluble or sparingly soluble bases – Ionic compounds that are insoluble in water,

consequently, not very basic.

• Formula Name

1. Cu(OH)2 copper (II) hydroxide

2. Fe(OH)2 iron (II) hydroxide

3. Fe(OH)3 iron (III) hydroxide

4. Zn(OH)2 zinc (II) hydroxide

5. Mg(OH)2 magnesium hydroxide

48

Aqueous Solutions: Weak Base

• Weak bases are covalent compounds that ionize slightly in water.

• Ammonia is most common weak base– NH3

• Many organic amines (Nitrogen compounds with C, H) are also weak bases.

– CH3NH2

– (CH3) 2NH

– (CH3) 3N

49


Bases• Weak bases are covalent compounds

that ionize slightly in water.

• Ammonia is the most common weak base

– NH3

Molecular NH3, and the ions NH4+ OH- are all

present in solution.

( ) ( ) ( )←→+ -

(aq)aq42g3 OH + NH OH + NHl

50

Aqueous Solutions: Solubility

Rules

Solubility Guidelines for Compounds in Aqueous Solutions–– It is very important that you know these guidelines It is very important that you know these guidelines

and how to apply them in reactions.and how to apply them in reactions.

1) Common inorganic acids and low-molecular-weight organic acids are water soluble.

2) All common compounds of the Group IA metal ions and the ammonium ion are water soluble.–– LiLi++, NaNa++, KK++, RbRb++, CsCs++, and NHNH44

++

51


Introduction3) Common nitrates, acetates, chlorates, and

perchlorates are water soluble.–– NONO33

--, CHCH33COOCOO--, ClOClO33--, and ClOClO44

--

4) Common chlorides are water soluble.– Exceptions – AgClAgCl, HgHg22ClCl22, & PbClPbCl22– Common bromides and iodides behave similarly to

chlorides.

– Common fluorides are water soluble.• Exceptions – MgFMgF22, CaFCaF22, SrFSrF22, BaFBaF22, and

PbFPbF22

52


Introduction5) Common sulfates are water soluble.

– Exceptions – PbSOPbSO44, BaSOBaSO44, & HgSOHgSO44– Moderately soluble – CaSOCaSO44, SrSOSrSO44, &

AgAg22SOSO44

6) Common metal hydroxides are water insolubleinsoluble.– Exceptions – LiOH, NaOH, KOH, RbOH LiOH, NaOH, KOH, RbOH &

CsOH (these are strong bases which are CsOH (these are strong bases which are all very soluble in water).all very soluble in water).

53


Introduction7) Common carbonates, phosphates, and

arsenates are water insolubleinsoluble.–– COCO33

22--, POPO4433--, & AsOAsO44

33--

– Exceptions- IA metals IA metals and NHNH44++

Ba(COBa(CO33))22 is moderately solubleis moderately soluble

– Moderately soluble – MgCOMgCO33

8) Common sulfides are water insolubleinsoluble.– Exceptions – IA metals IA metals and NHNH44

++ plus

IIA metalsIIA metals

54

Reactions in Aqueous Solutions

• Symbolic representation of what is

happening at the laboratory and molecular

levels in aqueous solutions.

– Copper reacting with silver nitrate.

• Laboratory level: The movie shows that

the Cu metal dissolves in the

Ag+ solution, while solid metallic

Ag appears.

55

Reactions in

Aqueous Solutions

• Symbolic representation of what is

happening at the laboratory and molecular

levels in aqueous solutions.

– Copper reacting with silver nitrate.

• Symbolic representation

( ) ( ) ( )s2(aq)3aq3s Ag 2)Cu(NOAgNO 2Cu +→+

56

Reactions in

Aqueous Solutions

• Another example of aqueous reactions.

– Sodium chloride reacting with silver nitrate.

• Laboratory level

• When AgNO3 and NaCl

aqueous solution are mixed

A white solid precipitate appears.

57

Reactions in

Aqueous Solutions

• Another example of aqueous reactions.

– Sodium chloride reacting with silver nitrate.

• Symbolic representation

( ) ( ) ( ) ( )aq3saqaq3 NaNOAgClNaCl AgNO +→+

58

Reactions in

Aqueous Solutions

There are three ways to write reactions in aqueous solutions.

1. Molecular equation – Show all reactants & products in molecular or ionic

form

2. Total ionic equation – Show the ions and molecules as they exist in solution

(s)(aq)4(aq)4(s) Cu + ZnSO CuSO + Zn →

( ) ( ) ( ) ( ) (s)

-2

aq4

2

aq

-2

aq4

2

aq(s) Cu +SO+ ZnSO+ Cu+Zn ++ →59

Reactions in

Aqueous Solutions

3. Net ionic equation

– Shows ions that participate in reaction and

removes spectator ions.

• Spectator ions do not participate in the

reaction.

60

Reactions in

Aqueous Solutions

• Look in total ionic equation for species that do

not change from reactant to product.

– Spectator ions in < >’s.

• Net ionic equation

( ) ( ) ( ) ( ) (s)

-2

aq4

2

aq

-2

aq4

2

aq(s) Cu +SO+ ZnSO+ Cu+Zn ++ →

( ) ( ) (s)

2

aq

2

aq(s) Cu + ZnCu + Zn ++ →

61

Reactions in

Aqueous Solutions

• In the total and net ionic equations the

only common substances that should be

written as ions are:

a. Strong acids

b. Strong bases

c. Soluble ionic salts

62

Oxidation Numbers

Guidelines for assigning oxidation numbers.1. The oxidation number of any free, uncombined element

is zero.

2. The oxidation number of an element in a simple (monatomic) ion is the charge on the ion.

Zero-valent elements: C(s), Ar(g), Cu(s), etc. and

1. The diatomic molecules H2(g), O2(g), N2(g), F2(g) , Cl2(g) , Br2(l) , I2(s)

2. Cu+2, Zn+2, Cl--, N-3, O-2 etc. 63

Oxidation Numbers

Guidelines for assigning oxidation numbers.3. In the formula for any compound, the sum of the

oxidation numbers of all elements in the compound is zero.

4. In a polyatomic ion, the sum of the oxidation numbers of the constituent elements is equal to the charge on the ion.

NaCl, CaO, Bal2 ; Mg(NO3)2, AlCl3, Na3PO4, Ca3(PO4) 2,

KCH3COO

Figure this out: What are the charges of the elements in

Ca(HSO4) 2 Fe2(HPO4)3 Zn(H2PO4)2

64

Oxidation Numbers for ions

5. Fluorine has an oxidation number of –1 in its compounds.

6. Hydrogen, H, has an oxidation number of +1 unless it is combined with metals, where it has the oxidation number -1.

– Examples – LiH, BaH2

7. Oxygen usually has the oxidation number -2.– Exceptions:

– In peroxides O has oxidation number of –1.

• Examples - H2O2, CaO2, Na2O2

– In OF2 ; O has oxidation number of +2.

65

Oxidation Numbers8. Use the periodic table to help with assigning

oxidation numbers of other elements.a. IA metals have oxidation numbers of +1.

b. IIA metals have oxidation numbers of +2.

c. IIIA metals have oxidation numbers of +3.

• There are a few rare exceptions.

d. VA elements have oxidation numbers of –3 in binarycompounds with H, metals or NH4

+.

e. VIA elements below O have oxidation numbers of –2 in binarycompounds with H, metals or NH4

+.

DRILLS

• NaNO3

• Ba(NO3)2

66

67

Oxidation Numbers

• Example 4-1: Assign oxidation numbers to each element in the following compounds:

• NaNO3

• Na = +1

• O = -2

• N = +5

– +1 + 3(-2) + x = 0

– x = +5

DRILLS

• K2Sn(OH)6

• You don’t have to memorize the

name of this specific

compound.

• It’s a coordination complex

called Potassium

tetrahydroxostannate ��68 69

Oxidation Number of Sn?

• K2Sn(OH)6• K = +1

• O = -2

• H = +1

• Sn = +5

– 2(+1) + 6(-2) + 6(+1) + x = 0

– x = +5 for Sn

70

Oxidation Number of Cl?

• HClO4

You do it!You do it!

• H = +1

• O = -2

• Cl = +7

71

Oxidation Number of N?

• NO2-

• O = -2

• N = +3

– 2(-2) + x = -1

– x = +3

72

Oxidation Number of C

• HCO3-

• O = -2

• H = +1

• C = +4

– +1 + 3(-2) + x = -1

– x = +4

73

Oxidation Numbers

• (COOH)2


• H = +1

• O = -2

• C = +3

• HOOC-COOH74

Naming Some

Inorganic Compounds• Binary compounds are made of two elements.

– metal + nonmetal = ionic compound

– nonmetal + nonmetal = covalent compound

• Name the more metallic element first.

– Use the element’s name.

• Name the less metallic element second.

– Add the suffix “ide” to the element’s stem.

75

Naming Some

Inorganic Compounds•• Nonmetal StemsNonmetal Stems

• Element Stem

• Boron bor

• Carbon carb

• Silicon silic

• Nitrogen nitr

• Phosphorus phosph

• Arsenic arsen

• Antimony antimon

76

Naming Some

Inorganic Compounds• Oxygen ox

• Sulfur sulf

• Selenium selen

• Tellurium tellur

• Phosphorus phosph

• Hydrogen hydr

77

Naming Some

Inorganic Compounds• Fluorine fluor

• Chlorine chlor

• Bromine brom

• Iodine iod

78

Naming Some

Inorganic Compounds

•• Binary Ionic Compounds Binary Ionic Compounds are made of a metal

cation and a nonmetal anion.

– Cation named first

– Anion named second

• LiBr lithium bromide

• MgCl2 magnesium chloride

• Li2S lithium sulfide

• Al2O3 You do it!You do it!

79

Naming Some

Inorganic Compounds• LiBr lithium bromide



• Al2O3 aluminum oxide

• Na3P You do it!You do it!

80

Naming Some

Inorganic Compounds• LiBr lithium bromide




• Na3P sodium phosphide

• Mg3N2 You do it!You do it!

81

Naming Some

Inorganic Compounds

• LiBr lithium bromide




• Na3P sodium phosphide

• Mg3N2 magnesium nitride

• Notice that binary ionic compounds with metals having

one oxidation state (representative metals) do not use

prefixes or Roman numerals.

82

Naming Some

Inorganic Compounds•• Binary ionic compounds containing Binary ionic compounds containing

metals that exhibit more than one metals that exhibit more than one

oxidation stateoxidation state

• Metals exhibiting multiple oxidation

states are:

1. most of the transition metals

2. metals in groups IIIA (except Al), IVA, & VA

83

Naming Some

Inorganic Compounds• There are two methods to name these

compounds.

1. Older method – add suffix “ic” to element’s Latin name for higher

oxidation state

– add suffix “ous” to element’s Latin name for lower oxidation state

2. Modern method

– use Roman numerals in parentheses to indicate metal’s oxidation state

84

Naming Some

Inorganic Compounds

• Compound Old System Modern System

• FeBr2 ferrous bromide iron(II) bromide

• FeBr3 ferric bromide iron(III) bromide

• SnO stannous oxide tin(II) oxide

• SnO2 stannic oxide tin(IV) oxide

• TiCl2 You do it!You do it!

85

Naming Some

Inorganic Compounds






• TiCl2 titanous chloride titanium(II) chloride


86

Naming Some

Inorganic Compounds







• TiCl3 titanic chloride titanium(III) chloride


87

Naming Some

Inorganic Compounds







• TiCl3 titanic chloride titanium(III) chloride

• TiCl4 does not workdoes not work titanium(IV) chloride

88

Naming Some

Inorganic Compounds

•• Pseudobinary ionic compoundsPseudobinary ionic compounds

• There are three polyatomic ions that commonly form binary ionic compounds.1. OH- hydroxide

2. CN- cyanide

3. NH4+ ammonium

• Use binary ionic compound naming system.

• KOH potassium hydroxide

• Ba(OH)2 barium hydroxide

• Al(OH)3 aluminum hydroxide

• Fe(OH)2 You do it!You do it!89

Naming Some

Inorganic Compounds• KOH potassium hydroxide



• Fe(OH)2 iron (II) hydroxide

• Fe(OH)3 You do it!You do it!

90

Naming Some





• Fe(OH)3 iron (III) hydroxide

• Ba(CN)2 You do it!You do it!

91

Naming Some






• Ba(CN)2 barium cyanide

• (NH4)2SYou do it!You do it!

92

Naming Some

Inorganic Compounds







• (NH4)2S ammonium sulfide

• NH4CN You do it!You do it!93

Naming Some

Inorganic Compounds







• (NH4)2S ammonium sulfide

• NH4CN ammonium cyanide

94

Naming Some

Inorganic Compounds•• Binary Acids Binary Acids are binary compounds consisting of hydrogen and a nonmetal.

• Compounds are usually gases at room temperature and pressure.– Nomenclature for the gaseous compounds is hydrogen (stem)ide.

• When the compounds are dissolved in water they form acidic solutions.– Nomenclature for the acidic solutions is

hydro (stem)ic acid.

95

Naming Some

Inorganic Compounds

• Formula Name Aqueous Solution

• HF hydrogen fluoride hydrofluoric acid

• HCl hydrogen chloride hydrochloric acid

• HBr hydrogen bromide hydrobromic acid

• H2S You do it!You do it!

96

Naming Some

Inorganic Compounds

• Formula Name Aqueous solution

• HF hydrogen fluoride hydrofluoric acid

• HCl hydrogen chloride hydrochloric acid

• HBr hydrogen bromide hydrobromic acid

• H2S hydrogen sulfide hydrosulfuric acid

97

Naming Some

Inorganic Compounds

•• Binary covalent molecular compounds Binary covalent molecular compounds composed of two nonmetals other than composed of two nonmetals other than hydrogenhydrogen

– Nomenclature must include prefixes that specify the number of atoms of each element in the compound.

• Use the minimum number of prefixes necessary to specify the compound.

– Frequently drop the prefix mono-.

98

Naming Some

Inorganic Compounds• Formula Name

• CO carbon monoxide

• CO2 carbon dioxide

• SO3 sulfur trioxide

• OF2 oxygen difluoride

• P4O6 You do it!You do it!

99

Naming Some Inorganic

Compounds• Formula Name





• P4O6 tetraphosphorus hexoxide

• P4O10 You do it!You do it!

100

Naming Some

Inorganic Compounds• Formula Name





• P4O6 tetraphosphorus hexoxide

• P4O10 tetraphosphorus decoxide

101

Naming Some

Inorganic Compounds

• The oxides of nitrogen illustrate why covalent compounds

need prefixes and ionic compounds do not.

• Formula Old Name Modern Name

• N2O nitrous oxide dinitrogen monoxide

• NO nitric oxide nitrogen monoxide

• N2O3 nitrogen trioxide dinitrogen trioxide

• NO2 nitrogen dioxide nitrogen dioxide

• N2O4 nitrogen tetroxide dinitrogen tetroxide

• N2O5 nitrogen pentoxide dinitrogen pentoxide

102

Naming Some

Inorganic Compounds

•• Ternary Acids and Their Salts Ternary Acids and Their Salts are made ofthree elements.– The elements are H, O, & a nonmetal.

• Two of the compounds are chosen as the basis for the nomenclature system.– Higher oxidation state for nonmetal is named (stem)ic acid.

– Lower oxidation state for nonmetal is named (stem)ous acid

• Salts are named based on the acids.– Anions of -ic acids make “ate” salts.

– Anions of -ous acids make “ite” salts.

103

Naming Some

Inorganic Compounds

• Names and Formulas of the Common “ic” acids– Naming these compounds will be easier if you have this list memorized.

• Group Name Formula

• IIIA boric acid H3BO3

• IVA carbonic acid H2CO3

silicic acid H4SiO4

• VA nitric acid HNO3

phosphoric acid H3PO4

arsenic acid H3AsO4104

Naming Some

Inorganic Compounds

• VIA sulfuric acid H2SO4

selenic acid H2SeO4

telluric acid H6TeO6

• VIIA chloric acid HClO3

bromic acid HBrO3

iodic acid HIO3

105

Naming Some

Inorganic Compounds• Salts are formed by the reaction of the acid with a strong base.

• Acid Salt

• HNO2 NaNO2

nitrous acid sodium nitrite

• HNO3 NaNO3

nitric acid sodium nitrate

• H2SO3 Na2SO3

sulfurous acid sodium sulfite

106

Naming Some

Inorganic Compounds• Acid Na Salt

• H2SO4 You do it!You do it!

107

Naming Some

Inorganic Compounds• Acid Na salt

• H2SO4 Na2SO4

sulfuric acid sodium sulfate

• HClO2 You do it!You do it!

108

Naming Some


• H2SO4 Na2SO4


• HClO2 NaClO2

chlorous acid sodium chlorite

• HClO3 You do it!You do it!

109

Naming Some


• H2SO4 Na2SO4


• HClO2 NaClO2


• HClO3 NaClO3

chloric acid sodium chlorate

110

Naming Some

Inorganic Compounds• There are two other possible acid and salt combinations.

• Acids that have a higher oxidation state than the “ic” acid are given the prefix “per”.– These acids and salts will have one more O atom than the “ic” acid.

• Acids that have a lower oxidation state than the “ous” acid are given the prefix “hypo”. – These acids and salts will have one less O atom than the “ous” acid.

111

Naming Some

Inorganic Compounds• Illustrate this series of acids and salts with the Cl ternary acids and salts.

• Acid Na Salt

• HClO NaClO

hypochlorous acid sodium hypochlorite

• HClO2 NaClO2


• HClO3 NaClO3

chloric acid sodium chlorate

• HClO4 NaClO4

perchloric acid sodium perchlorate

112

Naming Some

Inorganic Compounds

•• Acidic SaltsAcidic Salts are made from ternary acids that retain one or more of their acidic hydrogen atoms.– Made from acid base reactions where there is an insufficient amount of base to react with all of the hydrogen atoms.

• Old system used the prefix “bi”“bi” to denote the hydrogen atom.

• Modern system uses prefixes and the word hydrogen.

113

Naming Some

Inorganic Compounds

• NaHCO3

Old system sodium bicarbonate

Modern system sodium hydrogen carbonate

• KHSO4

Old system potassium bisulfate

Modern system potassium hydrogen sulfate

• KH2PO4

Old system potassium bis biphosphate

Modern system potassium dihydrogen phosphate

• K2HPO4 You do it!You do it!

114

Naming Some

Inorganic Compounds

• K2HPO4

Old system potassium biphosphate

Modern system potassium hydrogen phosphate

115

Naming Some

Inorganic Compounds•• Basic SaltsBasic Salts are analogous to acidic salts.

– The salts have one or more basic hydroxides

remaining in the compound.

• Basic salts are formed by acid-base reactions

with insufficient amounts of the acid to react with

all of the hydroxide ions.

• Use prefixes to indicate the number of hydroxide

groups.

116

Naming Some

Inorganic Compounds• Ca(OH)Cl

– calcium monohydroxy chloride

• Al(OH)Cl2– aluminum monohydroxy chloride

• Al(OH)2Cl You do it!You do it!

• aluminum dihydroxy chloride

There are many types ofchemical reactions.

We’ll name and study them now. ☺

REDOXExample of a

chemical reaction

119

Oxidation-Reduction Reactions:

An Introduction• Oxidation is an increase in the oxidation

number.

– Corresponds to the loss of electrons.

• Reduction is a decrease in the oxidation

number.

– Good mnemonic – reduction reduces the

oxidation number (high to low).

– Corresponds to the gain of electrons

120


An Introduction

• Oxidizing agents are chemical species that:1. oxidize some other substance

2. contain atoms that are reduced in the reaction

3. gain electrons

• Reducing agents are chemical species that:1. reduce some other substance

2. contain atoms that are oxidized in the reaction

3. lose electrons

REDOX reaction

• Corrosion of Iron:

• Fe(s) + O2 (g) � FeO, Fe2O3 (s) (rust),

• Actually the reaction is more complicated

than this. V

• Fe � Fe+2 + 2e-

• O2 + 2e- � O-2

• Fe � Fe+3 + 3e-121 122


An Introduction

• Two examples of oxidation-reduction or redox reactions.

• KMnO4 and Fe2+

– Fe2+ is oxidized to Fe3+

– MnO41- is reduced to Mn2+

• Combustion reactions are redox reactions

• Combustion of Mg– Mg is oxidized to MgO

– O2 is reduced to O2-

123


An Introduction

• Oxidation-reduction or redox reactions.

• Write the net redox reaction between Fe(NO3)3(aq) and Zn(s). The products are Fe(s) and Zn2+

• Write the net redox reaction between the reduction of bromine to bromide ions by Co(s).

The Co(s) is transformed into Co2+ ions.

124


An Introduction

• Example 4-2: Write and balance the formula unit,

total ionic, and net ionic equations for the oxidation

of sulfurous acid to sulfuric acid by oxygen in acidic

aqueous solution.

• Formula unit equation

• Total ionic equation


( ) ( ) ( )aq42g2aq32 SO H2 O SO H2 →+

( ) ( )−+ +→+ 2

4(aq)(aq)g2aq32 SO 2 H 4 O SOH 2 125


An Introduction



• Which species are oxidized and reduced?

• Identify the oxidizing and reducing agents.


( ) ( )−+ +→+ 2

4(aq)(aq)g2aq32 SO 2 H 4 O SOH 2

126


An Introduction

• H2SO3 is oxidized.

– The oxidation state of S in H2SO3 is +4.

– In SO42-, S has an oxidation state of +6.

• O2 is reduced.

– Oxidation state of O in O2 is 0

– In SO42-, O has an oxidation state of –2.

• H2SO3 is reducing agent.

• O2 is oxidizing agent.

127

Combination Reactions• Combination reactions occur when two or

more substances combine to form a compound.

• There are three basic types of combination reactions.1. Two elements react to form a new compound

2. An element and a compound react to form one new compound

3. Two compounds react to form one compound

128

Combination Reactions

1. Element + Element → CompoundA. Metal + Nonmetal → Binary Ionic Compound

( ) ( ) ( )sg2s NaCl 2ClNa 2 →+

129



( ) ( ) ( )sg2s MgO 2OMg 2 →+

130



( ) ( ) ( )s32s AlBr 2 Br3Al 2 →+l

131


1. Element + Element → CompoundB. Nonmetal + Nonmetal → Covalent Binary

Compound

( ) ( ) ( )s104g2s4 O PO 5P →+

132


1. Element + Element → CompoundB. Nonmetal + Nonmetal → Covalent Binary

Compound

( ) ( ) ( )l3g2s4 PCl4 Cl 6P →+

133


1. Element + Element → Compound

B. Nonmetal + Nonmetal → Covalent Binary

Compound

• Can control which product is made with the

reaction conditions.

( ) ( ) ( )

chlorine limitedin

AsCl 2 Cl 3As 2 s3g2s →+

( ) ( ) ( )

chlorine excessin

AsCl 2 Cl 5As 2 s5g2s →+134


1. Element + Element → Compound

B. Nonmetal + Nonmetal → Covalent Binary

Compound

• Can control which product is made with the

reaction conditions.

( ) ( ) ( )

fluorine limitedin

SeF F 2Se s4g2s →+

( ) ( ) ( )

fluorine excessin

SeF F 3Se g6g2s →+135


2. Compound + Element → Compound

( ) ( ) ( )s5g2s3 AsClClAsCl →+

( ) ( ) ( )g6g2s4 SFFSF →+

136


The reaction of oxygen with oxides of

nonmetals is an example of this type of

combination reaction.

( ) ( ) ( )g3

&catalyst

g2g2 SO 2OSO 2 →+∆

( ) ( ) ( )g2g2g CO 2OCO 2 →+

104264 OPO 2OP →+137


3. Compound + Compound → Compound

– gaseous ammonia and hydrogen chloride

– lithium oxide and sulfur dioxide

( ) ( ) ( )s4gg3 ClNH HClNH →+

3222 SO LiSOOLi →+

138

Decomposition Reactions

• Decomposition reactions occur when one

compound decomposes to form:

1. Two elements

2. One or more elements and one or more

compounds

3. Two or more compounds

139


1. Compound → Element + Element

• decomposition of dinitrogen oxide

• decomposition of calcium chloride

( ) ( ) ( )g2g2g2 ON 2ON 2 +→∆

( ) ( ) ( )g2

yelectricit

2 ClCaCaCl + →ll

( ) ( ) ( )l2s

h

s BAg 2AgBr 2 r+→ ν

• decomposition of silver halides

140


2. Compound → One Element +

Compound(s)

– decomposition of hydrogen peroxide

( ) ( ) ( )g22

or Mn or Feνh

aq22 OO H2O H23

+ →+

l

141


3. Compound → Compound + Compound

– decomposition of ammonium hydrogen carbonate

( ) ( ) ( ) ( )g2g2g3s34 COO HNHHCONH ++→∆

142

Displacement Reactions

•• Displacement reactionsDisplacement reactions occur when one

element displaces another element from a

compound.

– These are redox reactions in which the more

active metal displaces the less active metal of

hydrogen from a compound in aqueous

solution.

– Activity series is given in Table 4-14.

143


1. [More Active Metal + Salt of Less Active Metal] → [Less

Active Metal + Salt of More Active Metal]

– molecular equation

( ) ( ) (s)aq3(s) aq3 Ag CuNO Cu +AgNO +→

144






( ) ( ) ( ) ( ) ( ) (s)

-

aq3aqs

-

aq3aq Ag NO+CuCu +NO+ Ag +→++

( ) ( ) (s)aq(s)aq Ag Cu Cu +Ag +→ ++

145


2. [Active Metal + Nonoxidizing Acid] → [Hydrogen +

Salt of Acid]

– Common method for preparing hydrogen in the laboratory.

– HNO3 is an oxidizing acid.

• Molecular equation

( ) ( ) ( )g2aq342aq42(s) H 3 + )(SOAl SO3H + Al 2 →

146






( ) ( ) ( ) ( ) ( )g2

-2

aq4

3

aq

-2

aq4aq(s) H 3 + SO 3 + Al 2 SO 3+H 6 + Al 2 ++ →

( ) ( ) ( )g2

3

aqaq(s) H 3 +Al 2 H 6 + Al 2 ++ →

147


• The following metals are active enough to

displace hydrogen

– K, Ca, Na, Mg, Al, Zn, Fe, Sn, & Pb

• Notice how the reaction changes with an

oxidizing acid.

– Reaction of Cu with HNO3.

• H2 is no longer produced.

148


3. [Active Nonmetal + Salt of Less Active Nonmetal] → [Less

Active Nonmetal + Salt of More Active

Nonmetal]


( ) ( ) ( ) (aq)s2aqg2 NaCl 2 I NaI 2 + Cl +→



( ) ( ) ( ) ( ) ( ) ( )-

aqaqs2

-

aqaqg2 Cl 2 +Na 2 I I 2 + Na 2 +Cl ++ +→149




( ) ( ) ( ) ( )-

aqs2

-

aqg2 Cl 2 I I 2 +Cl +→

150

Metathesis Reactions•• Metathesis reactionsMetathesis reactions occur when two ionic

aqueous solutions are mixed and the ions switch partners.

AX + BY → AY + BX

• Metathesis reactions remove ions from solution in two ways:

1. form predominantly unionized molecules like H2O

2. form an insoluble solid

• Ion removal is the driving force of metathesis reactions.

151

Metathesis Reactions

1. Acid-Base (neutralization) Reactions

– Formation of the nonelectrolyte H2O

– acid + base → salt + water

152



)(2 (aq)(aq)(aq) OH + KBr KOH + HBrl

→

�Total ionic equation


( ) ( ) ( ) ( ) ( ) ( ) )(2

-

aqaq

-

aqaq

-

aqaq OH + Br+KOH+K+Br+Hl

+++ →

�Net ionic equation


( ) ( ) )(2

-

aqaq OH OH +Hl

→+

153



)(2aq)(23(aq)3(aq)2 OH 2 + )Ca(NOHNO 2 + Ca(OH)l

→

�Total ionic equation


( ) ( ) ( ) ( ) ( ) ( ) )(2

-

aq3

2

aq

-

aq3aq

-

aq

2

aq OH 2 +NO 2+ CaNO 2+ H 2+OH 2+Cal

+++ →

�Net ionic equation


( ) ( )

( ) ( ) )(2aq

-

aq

)(2aq

-

aq

OH H+OH

betteror

OH 2 H 2+OH 2

l

l

→

→

+

+

154


2.2. Precipitation reactionsPrecipitation reactions are metathesis reactions in which an insoluble compound is formed.

– The solid precipitates out of the solution much like rain or snow precipitates out of the air.

155


• Precipitation Reactions


(s)3)aq(3aq)(32(aq)23 CaCO +KNO 2 COK + )Ca(NO →

�Total ionic reaction


( ) ( ) ( ) ( )

( ) ( ) ( )s3

-

aq3aq

-2

aq3aq

-

aq3

2

aq

CaCO NO 2K 2

COK 2 NO 2 Ca

++

→+++

+

++

156


• Net ionic reaction


( ) ( ) (s)3

-2

aq3

2

aq CaCO CO +Ca →+

157



( )2(s)43)aq(aq)(43(aq)2 POCa +NaCl 6 PONa 2 + CaCl 3 →



( ) ( ) ( ) ( )

( ) ( ) ( ) ( )s243

-1

aq

1

aq

-3

aq4

1

aq

-1

aq

2

aq

POCa +Cl 6 Na 6

PO2 Na 6 + Cl 6 Ca 3

+

→++

+

++

158




( ) ( ) ( ) ( )s243

-3

aq4

2

aq POCa PO 2 Ca 3 →++

159



( ) ( )g22)aq(aq)(32(aq) SO O H+NaCl 2 SONa + HCl2 +→l



( ) ( ) ( ) ( )

( ) ( ) ( ) ( )g22

-1

aq

1

aq

-2

aq3

1

aq

-1

aq

1

aq

SO OH +Cl 2Na 2

SO Na 2 + Cl 2H 2

++

→++

+

++

l

160




161

Gas-Formation Reactions

• A gas-formation reaction is a type of

reaction in which there is a formation of an

insoluble or slightly soluble gas when

there are no gaseous reactants.

• Displacement reactions in which an active

metal displaces from an acid or from water

are gas-formation reactons; they are not

methathesis reactions.

162


• Consider hydrochloric acid with calcium

carbonate to form carbonic acid.

Formula Unit

2HCl(aq) + CaCO3 (s) → H2CO3 (aq) + CaCl2 (aq)

Total Ionic

2[H+(aq) + Cl-(aq)] → H2CO3 (aq)

+ CaCO3 (s) + [Ca2+(aq) + 2Cl-(aq)]

Net Ionic

2H+(aq) + CaCO3 (s) → H2CO3 (aq) + Ca2+

(aq)

163


• Enough heat is generated in the reaction

to cause thermal decomposition of

carbonic acid.

H2CO3 (aq) → CO2 + H2O(l)

• The net effect of the chemical reqaction

and subsequent decomposition is

2HCl(aq) + CaCO3 (s) → CO2 + H2O(l) + CaCl2 (aq)

164

Synthesis Question

• Barium sulfate is a commonly used imaging

agent for gastrointestinal X-rays. This

compound can be prepared by some of the

simple reactions described in this chapter. Write

a balanced aqueous reaction for the production

of barium sulfate. You can choose any aqueous

starting materials that will form barium sulfate!

165

Synthesis Question

• Find two aqueous soluble compounds that

have Ba in one compound and SO42- in the

second. When they are mixed, the barium

sulfate will precipitate out. One possibility

is:

)s(4(aq))aq(422(aq) BaSO + NaCl 2 SONa +BaCl →

Practice Exercises

Identify the type of reaction involved. Check

also if they are redox reactions. Write the net

ionic equation.

1. Al(s) + FeBr3 (aq) � AlBr3 (aq) + Fe(s)

2. NaBr(aq)+AuNO3(aq) �AuBr(s)+NaNO3(aq)

3. Mg(s) + H2O (l) � MgO(s) + H2(g)

4. BaCO3(s) � BaO(s) + CO2(g)

5. Zn(s) + 2HNO3(aq) � Zn(NO3 )2(aq)+ H2(g)

6. PCl3(g) + Cl2(g) � PCl5(g) 166

Practice Exercises:

Concentrations of Solution, M.1. What is the molarity of a solution prepared by

dissolving 355 g of sodium phosphate (MM: 163.94) in

water and diluting into 4.50 L of solution?

2. How many moles of NaCl are present in a 850 mL of

0.50 M NaCl(aq) solution?

3. Calculate the final volume of a solution obtained in

100mL of 12.0 M NaOH is diluted to make it 5.20 M in

concentration.

4. What is the molarity of a solution prepared by mixing

35.0 mL of 0.375 M NaCl with 47.5 mL of 0.632 M

NaCl?

167

Pratice in Percent by Mass/Mixed

1. The density of 18.0% solution of ammonium sulfate is

1.10g/mL. What mass of (NH4)2SO4 (MM=132.14) is

required to prepare a 775.0 mL of this solution?

2. How many moles and grams of solute are contained in

750.0 g of a 15.00% aqueous solution of K2Cr2O7?

MM K2Cr2O7 : 294.19

3. What mass of AgCl could be formed by mixing 10.0 mL of

1.20% NaCl by mass solution (d=1.02g/mL) with 50.0 mL

of 1.21x10-2 M AgNO3? The is metathesis forming AgCl(s).

168

Practice Exercises: Percent

Yield• PCl3(g) + Cl2(g) � PCl5(g)

• Suppose the percent yield for the reaction is 86.5%.

What mass of PCl5(g) is obtained from the reaction of

96.7g with excess Cl2(g)?

• Molar Mass(g/mol):PCl5:208.27; PCl3:137.35

169

• CH4(g) + 4S(g) � CS2 (g) + H2S (g)

• The percent yield for the reaction is consistently 87.0%.

How many grams of sulfur would be needed to obtain

80.0 g of CS2(g).

• Molar Mass(g/mol):CS2:76.14; S:32.06

Practice exercise: Mixed

• An iron ore that contains Fe3O4 reacts according to the

reaction Fe3O4 (s) + 2C (s) � 3Fe(s) + 2CO2(g)

• We obtain 3.49 g of Fe from the reaction of 75.0 g of the

ore with C(s) in a furnace.

(a) How many moles and grams of Fe3O4 are present in the

ore?

(b) What is the percent Fe3O4 in the ore?

(c) How many moles and grams of C reacted with the ore?

(d) In 10.0 grams of Fe3O4, calculate the number of Fe and O

atoms.

(d2) Calculate the mass of O (in grams) in 10.0 g Fe3O4

170

Practice Exercises: Mixed

1. Calculate the volume of 2.25 M phosphoric acid

solution necessary to react with 45.0 mL of 0.150 M

Mg(OH)2? The reaction is metathesis (acid-base) rexn.

2. Magnesium oxide, marketed as milk of magnesia, is a

common antacid. What volume in mL of gastric juice

(HCl) corresponding to acidity of 0.17 M HCl could be

neutralized by 104. mg of MgO? The rexn is

metathesis. MM MgO:40.31 g/mol.

3. Consider the decomposition of a 15.0g KClO3

[(MM:103.1g/mol) which is 76.5% by weight KClO3] into

KCl(s) and O2(g), calculate the number of moles and

grams O2(g) produced. How many atoms of O are

produced? 171

More on identifying reactions.

1. Sulfuric acid + 2KOH � ? + 2water.

2. 2Rb(s) + Br2(l) � ?RbBr(s).

3. 2KI + fluorine (g) � 2KF(aq) + I2(s)

4. CaO(s) + Si-dioxide(s) � CaSiO3(s)

5. S(s) + oxygen(g) � SO2(g)

6. HgS(s) + oxygen (g) � Hg(l) + SO2(g)

7. Pb(s) + 2HBr(aq) � PbBr2(s) + H2(g).

8. N2O5(s) + H2O(l) � ?HNO3(aq)172

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