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Compounds & Chemical Formulas. Common Compounds. H 2 O water the most common compound on the earth’s surface SiO 4 silicate the most common compound in the earth’s crust. There are millions of compounds!. Common Compounds. C 6 H 12 O 6 glucose C 55 H 72 O 5 N 4 Mg chlorophyll. H. O.

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slide2

Common Compounds

  • H2O
    • water
    • the most common compound on the earth’s surface
  • SiO4
    • silicate
    • the most common compound in the earth’s crust
slide3

There are millions of compounds!

Common Compounds

  • C6H12O6
    • glucose
  • C55H72O5N4Mg
    • chlorophyll
slide4

H

O

H

C

H

O

O

C

H

H

O

C

C

H

O

C

H

O

H

slide5

matter

puresubstances

mixtures

slide6

matter

puresubstances

compounds

mixtures

elements

iron

gold

lead

uranium

sodium chloride

glucose

calcium carbonate

water

slide7

matter

puresubstances

mixtures

elements

compounds

iron

gold

lead

uranium

sodium chloride

glucose

calcium carbonate

water

slide8

homogeneousmixtures

heterogeneousmixtures

mixtures

air

gasoline

14 karat gold

seawater

granite

soil

blood

chocolate cake

slide9

matter

puresubstances

mixtures

elements

compounds

homogeneousmixtures

heterogeneousmixtures

iron

gold

lead

uranium

sodium chloride

glucose

calcium carbonate

water

air

gasoline

14 karat gold

seawater

granite

soil

blood

chocolate cake

slide10

PureSubstances

contain only one kind of matter

slide11

Elements

All atoms are the same.

slide12

Compounds

  • made up entirely of the same molecules or ions
  • have the same ratio of elements
slide13

Compounds

2 types

Organic

Inorganic

  • do not have carbon
  • water, table salt
  • contain carbon
  • living things
slide14

Mixtures

  • combinations of substances
  • physically placed together, not bonded
  • no set ratio
slide15

Mixtures

2 types

Homogeneous

Heterogeneous

slide16

Homogeneous

  • Particles about the same size
  • Separate types of particles not usually visible
  • Called solutions

Examples: saltwater, alloys

slide17

Heterogeneous

  • nonuniform—different sized particles
  • Separate types of particles can often be seen.

Examples: soil, vegetable stew, wood, concrete

slide18

Chemical Formulas

  • A shorthand way to show the composition of a pure substance
  • Show how many atoms of each element make up a substance
slide19

Chemical Formulas

  • Example: Rust
  • Its chemical formula is Fe2O3, iron oxide.
  • The 2 and 3 are subscripts which tell how many atoms of each element are present in one molecule of rust.
slide20

What part of a chemical formula tells the number of atoms of each element?

  • Subscript
  • Molecule
  • Empirical formula
  • Definite composition
slide21

Chemical Formulas

  • Empirical Formula
  • Gives only the ratio of atoms
  • Not necessarily unique to a substance
  • May be the same as the actual chemical formula
slide22

We can predict the formula based on the number of valence electrons.

Chemical Formulas

  • How do we know that the formula for table salt is NaCl?
  • Why can’t it be NaCl2 (as in MgCl2,magnesium chloride)?
slide23

Oxidation Number

The number of electrons that a bonded atom would have to gain or lose to return to its neutral state.

slide24

Oxidation Number

  • If a metal has lost electrons, it will have to regain them to return to its neutral state.
    • Therefore, it has a positive oxidation number.
  • A nonmetal tends to gain electrons, so it will have a negative oxidation number.
slide25

+ 1

+ 2

Li

Be

+ 3

+ 4

B

C

– 4

slide26

– 3

– 2

N

O

– 1

0

F

Ne

slide27

Oxidation Number

  • How do we know the oxidation number of an atom?
    • By knowing which group it is in.
slide28

Oxidation Number

  • Group 1A (alkali) metals have one outer electron, which they tend to lose.
    • What would be the charge of the ion if it loses one electron?

+ 1

slide29

Oxidation NumberExample

  • Sodium

11 electrons (–) 11 protons (+)

  • If it loses one electron:
  • Sodium

10 electrons (–) 11 protons (+)

= 0

= +1

slide31

Oxidation Number

  • Group 2A metals (alkaline-earth) lose two electrons.
    • What oxidation number do they have?

+ 2

slide32

Oxidation Number

  • What is the oxidation number of
    • Group 3A elements?
    • Group 8A elements?
    • Group 7A elements?
    • Group 6A elements?

+ 3

0

– 1

– 2

slide33

Oxidation Number

  • What is the oxidation number of
    • Group 5A elements?
    • Group 4A elements?

+ 5, – 3

+ 4, – 4

slide34

To write a chemical formula, we must know the

  • atoms.
  • molecules.
  • oxidation numbers.
  • substance.
slide35

What tells us the oxidation number of an atom?

  • Its size
  • Its group
  • Its color
  • Its subscripts
slide38

0

0

0

0

Na Ar O2 S8

Oxidation Number Rules

Rule 1: Free Element Rule

The oxidation number (ON) of pure elements is 0.

slide39

2+

3+

4+

+

K Fe AlSnCl

Oxidation Number Rules

Rule 2: Ion Rule

The ON of a monatomic ion equals its charge.

+1

+2

+3

+4

–1

slide40

Oxidation Number Rules

  • Rule 3: Specific ON Rule
  • Certain families have the same ON.
    • Alkali metals = +1
    • Alkaline earth metals = +2
slide41

Oxidation Number Rules

  • Rule 3: Specific ON Rule
  • Certain families have the same ON.
    • Hydrogen = +1 (when bonded to another nonmetal)
    • Hydrogen = –1 (when bonded to metals)
slide42

Oxidation Number Rules

  • Rule 3: Specific ON Rule
  • Certain families have the same ON.
    • Oxygen = –2 (except when bonded to fluorine)
    • Halogens = –1 (when bonded to metals)
slide43

Oxidation Number Rules

Rule 4: Zero Sum Rule

The sum of the oxidation numbers of all atoms in a compound must equal zero.

+1

–1

0

+

Na + ClNaCl

slide44

Oxidation Number Rules

Rule 4: Zero Sum Rule

The sum of the oxidation numbers of all atoms in a compound must equal zero.

0

0

(+1)

(+1)

2

–2

-3

3

H2O

NH3

slide45

Oxidation Number Rules

Rule 4: Zero Sum Rule

The sum of the oxidation numbers of all atoms in a compound must equal zero.

0

0

(–1)

+3

3

(–1)

+2

2

AlCl3

CaI2

slide48

Writing Chemical Formulas

If we know the elements and their oxidation numbers, we can predict the formula of a compound.

slide49

Writing Chemical Formulas

  • Write the less electronegative element first.
  • Find the oxidation number of the elements.

+3

–1

Al

F

slide50

Writing Chemical Formulas

  • Change the number of atoms of the elements so that the sum of the oxidation numbers equals 0.

0

+3

–1

(–1)

3

F3

Al

F

slide51

Writing Chemical Formulas

  • Change the number of atoms of the elements so that the sum of the oxidation numbers equals 0.

0

(+2)

2

(–3)

3

Be

Be3

N

N2

slide52

Polyatomic Ions

A polyatomic ion is two or more covalently-bonded atoms that act as a single particle which gains or loses electrons.

slide53

Oxidation Number Rules

Rule 5: Polyatomic Ion Charge Rule

The sum of the oxidation numbers of all atoms in a polyatomic ion must equal its charge.

slide54

Oxidation Number Rules

Rule 5: Polyatomic Ion Charge Rule

Example: NO3

N

1(+5)

= +5

O

3(–2)

= –6

–1

slide55

Oxidation Number Rules

Rule 5: Polyatomic Ion Charge Rule

Example: PO4

–3

P

1(+5)

= +5

O

4(–2)

= –8

–3

slide56

Write the formula for the product of the reaction between beryllium and oxygen.

  • BeO2
  • Be2O
  • BeO
  • BeO3
slide57

Write the formula for calcium hydroxide.

  • Ca(OH)
  • Ca(OH)2
  • Ca(OH)3
  • Ca2(OH)
slide58

Naming Compounds

There are three types of compounds that we will name.

Binary ionic

Binary covalent

Polyatomic ionic

slide59

Naming Compounds

  • Binary Ionic Compounds
  • Have only two elements: a metal and a nonmetal
  • When naming them:
    • Name the metal first.
    • Change the nonmetal ending to “-ide.”
slide60

NaF

sodium fluoride

K2S

potassium sulfide

Ag2O

silver oxide

zinc nitride

Zn3N2

AlBr3

aluminum bromide

slide61

Naming Compounds

  • Binary Ionic Compounds
  • If the metal has more than one ON, use the Stock system.
    • A Roman numeral, in parenthesis, comes after the metal name.
    • This number indicates the ON of the metal.
slide62

Naming Compounds

  • Binary Ionic Compounds
  • The Stock system is used with iron, copper, mercury, tin, and lead.
slide63

Naming Compounds

  • Binary Covalent Compounds
  • A prefix is placed before one or both of the elements’ names to indicate how many atoms of each element are in the molecule.
slide64

Name PbCl2.

  • Lead chloride
  • Lead dichloride
  • Lead (III) chloride
  • Lead (II) chloride
slide65

Naming Compounds

  • Polyatomic Ions
  • Compounds with polyatomic ions are named the same way as binary ionic compounds.
slide66

Naming Compounds

  • Polyatomic Ions
  • The first name is the metal’s name.
    • Exception: NH4 can come first.
  • The second name is the polyatomic ion.

+

slide67

NH4Cl

ammonium chloride

NaOH

sodium hydroxide

K2CO3

potassium carbonate

ammonium phosphate

(NH4)3PO4

BaSO4

barium sulfate

slide68

Naming Compounds

  • Common names, such as soda ash or Epsom salts, do not give any information about the chemical formula and are not usually used.
slide69

Name Al2(SO4)3.

  • Aluminum sulfite
  • Aluminum sulfate
  • Ammonium sulfite
  • Ammonium sulfate
slide71

Chemical Change

Occurs when atoms rearrange to form new substances

slide72

Chemical Change

  • In a chemical reaction, atoms are not destroyed.
  • The original bonds in the substances are broken, and new bonds form to produce different substances.
slide73

Chemical Change

  • Signs that a chemical reaction MAY have occurred
  • Solid separates
  • Gas produced
  • Color changes
  • Temperature changes
  • Light, sound, or other energy produced
slide74

Chemical Change

  • However, these changes are only signs; the guarantee that a chemical reaction has occurred is the formation of new substances.
slide75

What is the only guarantee that a chemical reaction has occurred?

  • Heat is produced.
  • Light is produced.
  • A new substance is formed.
  • Solid separates from a liquid.
slide76

Chemical Equation

Describes the process of a chemical reaction using symbols

slide77

Chemical Equations

  • Because of the law of conservation of matter, the number of atoms of each element must be the same on both sides of the equation.
slide78

Example

Zinc reacts with oxygen. Describe the process in a chemical equation.

Zn + O2

slide79

Example

Zinc reacts with oxygen. Describe the process in a chemical equation.

0

+2

–2

ZnO

slide80

Example

Zinc reacts with oxygen. Describe the process in a chemical equation.

Zn + O2ZnO

product

reactants

slide81

reactants

products

C

+

D

A

+

B

slide82

Chemical Equations

  • Did you notice a problem with

Zn + O2ZnO?

The number of oxygen atoms is not the same on both sides, so the equation is not balanced.

slide83

Chemical Equations

  • Chemical equations must be balanced because of the law of conservation.
  • To balance an equation, coefficients are placed in front of the formulas as needed.

2

2

Zn + O2ZnO

slide84

Example

3

4

2

Al + O2

Al2O3

slide85

Example

2

2

H2 + O2

H2O

slide86

Balance the equation.

CH4 C + H2

  • 2 CH4 2 C + 3 H2
  • CH4 C + 2 H2
  • 3 CH4 3 C + 4 H2
  • 2 CH4 2 C + 5 H2
slide87

Equation Symbols

Symbol

Meaning

slide88

Equation Symbols

Symbol

Meaning

slide89

Equation Symbols

Symbol

Meaning

slide90

Chemical Equations

  • Catalysts are substances which are involved in the reaction but do not become a new substance.
    • They usually speed up the reaction.
    • They are written above the arrow.
slide92

Types of Reactions

  • There are four types of reactions.
  • Combination
  • Decomposition
  • Single replacement
  • Double replacement
slide93

Types of Reactions

  • Composition Reactions
  • Occur when two substances unite
  • Also called synthesis or combination reactions
  • Always have a single product

X+Y

XY

slide94

Types of Reactions

  • Decomposition Reactions
  • Occur when a substance splits into parts
  • Always have a single reactant

XY

X +Y

slide95

Types of Reactions

  • Single Replacement Reactions
  • Occur when a more active substance replaces one part of a compound that is less active

+

+

X

X

Y

Y

Z

Z

slide96

Types of Reactions

  • Double Replacement Reactions
  • Occur when parts of two substances exchange
  • Often make a precipitate

+

+

W

W

X

X

Y

Y

Z

Z

slide97

Of the following reactions, which one is a single replacement reaction?

  • 2BaO2 2BaO + O2
  • Mg(OH)2 + 2HCl MgCl2 + 2H2O
  • SiCl4 + 2H2 Si + 4HCl
  • N2 + 3H2 2NH3
slide98

Of the following reactions, which one is a double replacement reaction?

  • 2BaO2 2BaO + O2
  • SiCl4 + 2H2 Si + 4HCl
  • N2 + 3H2 2NH3
  • Mg(OH)2 + 2HCl MgCl2 + 2H2O
slide99

Of the following reactions, which one is a composition reaction?

  • 2BaO2 2BaO + O2
  • Mg(OH)2 + 2HCl MgCl2 + 2H2O
  • SiCl4 + 2H2 Si + 4HCl
  • N2 + 3H2 2NH3
slide100

Of the following reactions, which one is a decomposition reaction?

  • Mg(OH)2 + 2HCl MgCl2 + 2H2O
  • 2BaO2 2BaO + O2
  • SiCl4 + 2H2 Si + 4HCl
  • N2 + 3H2 2NH3
slide101

Exothermic Reactions

  • Exothermic reactions are reactions that give off heat; heat is a product.
  • Burning (combustion) is an exothermic reaction.
  • These reactions produce heat, so they keep themselves going.
slide102

Endothermic Reactions

  • Endothermic reactions are reactions that absorb heat; heat is considered to be a reactant.
  • Most decomposition reactions are endothermic.
  • Since they absorb heat, these reactions tend to stop when a source of heat is removed.