ppt inorganic chemistry

Chapter 2: Compounds and

Chemical Reactions

• Essentially all elements combine to form

compounds

• Compounds are of two types:

– Molecular, which involve shared electrons and

consist of electrically neutral, discrete particles

called

molecules

•

Chemical formulas

are collections of

chemical symbols that are used to describe

elements and compounds

– Free elements are not combined with other

elements in a compound

• Examples: Fe (iron), Na (sodium), and K (potassium)

• Chemical formulas

specify

the composition

of a substance

• NaCl is composed of the elements sodium and chlorine in a one-to-one (atom) ratio

• Fe2O3 is composed of the elements iron and oxygen

in a two-to-three ratio

• CO(NH2)2 expands to CON2H4, but there are good

reasons to write some compounds with parentheses • Hydrates are crystals that contain water molecules,

for example plaster: CaSO4 •2H2O

•

Chemical equations

describe what happens

in a chemical reactions

• Hydrogen and oxygen combine to form

water

• Hydrogen and oxygen are called reactants

• Water is called the product

• Reactants are separated from products with “”

2 H

₂

+ O

₂

 2 H

₂

O

• Note that the “” is like an equal sign because both

• This can be represented as:

Note: Mass is conserved because the number of atoms of each type remains the same on each side of the arrow.

(Both sides of the arrow show 4 H and 2 O atoms.) This equation is said to be

balanced.

The “2” in front of formulas H₂ and H₂O are called

coefficients. They indicate the number of molecules of each type and can change when balancing a chemical equation.

• It is sometimes useful to include the

physical state of reactants and products

• For solids use s, liquids use l, gases use g, and for aqueous solutions use aq.

• For example, the reaction between stomach

acid (an aqueous solution of HCl) and

sodium carbonate (an antacid) can be

written

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

• Almost all chemical reactions either absorb

or give off energy, often as

heat

or

light

•

Kinetic

and

potential

energy are both

important in chemistry

– Kinetic energy is the energy an object has when

moving

– Potential energy is the energy an object has due

to its position

• Energy must also be conserved

– The Law of Conservation of Energy:

• Energy cannot be created or destroyed; it can only be converted from one form to another

•

Heat

and

temperature

are related to kinetic

energy

• The temperature of an object is proportional to its average kinetic energy (average speed of its atoms) • Heat or thermal energy is transferred between

objects with different temperatures

• Chemical energy is a form of potential

energy

• The analysis of temperature changes in

chemical reactions can provide information

about the potential energy changes that

occur

– The kinetic molecular theory of matter

provides more details about chemical energy

changes and is discussed in Chapter 7

• As a general rule, molecular compounds are

formed when nonmetallic elements

combine

• Many molecular compounds contain

hydrogen:

Group Noble Period IVA VA VIA VIIA Gas 2 CH₄ NH₃ H₂O HF Ne 3 SiH₄ PH₃ H₂S HCl Ar 4 GeH₄ AsH₃ H₂Se HBr Kr 5 SbH₃ H₂Te HI Xe

•

Organic chemistry

is a major specialty that

deals with compounds containing mostly

carbon and hydrogen

•

Hydrocarbons

contain only hydrogen and

carbon and are organic compounds

•

Alkanes

are the simplest hydrocarbons

– General formula is C

n

H

2n+2

• Other classes of hydrocarbons exist

• Different classes of organic compounds are

derived

from hydrocarbons by replacing

hydrogen

• For example

alcohols

result when a

H

is

replaced by

OH

in a hydrocarbon

Methanol (wood alcohol),

•

Inorganic compounds

are substances

not

considered to be derived from hydrocarbons

• The rules for naming, or

nomenclature

, of

simple inorganic compound is covered now

(organic nomenclature is covered later)

•

Binary compounds

are compounds

comprised of two different elements

• The first element in the formula is identified

by its English name, the second by

appending the suffix

–ide

to its

stem

Chemical Name as Name as

Symbol Stem First Element Second Element O ox- oxygen oxide

• The number of each type of atom is

specified with Greek prefixes

Greek Prefixes

mono- = 1 (often omitted) hexa- = 6 di- = 2 hepta- = 7 tri- = 3 octa- = 8 tetra- = 4 nona- = 9 penta- = 5 deca- = 10

Examples:

PF₅ = phosphorus pentafluoride HCl = hydrogen chloride

N₂O₅ = dinitrogen tetraoxide

Note: many

compounds have

• The subscripts in the formula of an ionic

compound always specifies the smallest

whole-number ratio of the ions because

molecules don’t exist in ionic compounds

• The smallest unit of a compound is called

the

formula unit

• Positively charged ions have more protons than electrons and are called cations

• Negatively charged ions have more electrons than protons and are called anions

• Ionic compounds are composed of charged

particles (ions)

– Ions can be formed from the reaction of metal

with a nonmetal

• The charges on many representative

elements can be predicted:

– Metals form cations

• The positive charge on the cation is the same as the “A” group number of the metal

– Nonmetals form anions

• The negative charge on the anion is equal to the number of spaces to the right we have to move in the periodic table to get to a noble

• Rules for writing Formulas of Ionic

Compounds:

1) The positive ion is given first in the formula.

2) The subscripts in the formula must produce an

electrically neutral formula unit.

3) The subscripts should be the set of smallest

whole numbers possible.

• Ions formed by transition metals (Group

IIIB – VIIIB) and post-transition metals:

Transition Metals

Chromium Cr2+, Cr3+ Zinc Zn2+ Manganese Mn2+, Mn3+ Silver Ag+ Iron Fe2+, Fe3+ Cadmium Cd2+

Cobalt Co2+, Co3+ Gold Au+, Au3+ Nickel Ni2+ Mercury Hg₂2+, Hg2+ Copper Cu+, Cu2+

Post-transition Metals

• Some

polyatomic

ions (ions with two or

more atoms):

Ion

Name

Ion

Name

NH

₄+

ammonium ion CO

₃2-

carbonate ion

OH

-

hydroxide ion H

3

O

+

hydronium ion

NO

2-

nitrite ion

SO

32-

sulfite ion

NO

₃-

nitrate ion

SO

₄2-

sulfate ion

ClO

₂-

chlorite ion

CrO

₄2-

chromate ion

ClO

3-

chlorate ion

Cr

2

O

72-

dichromate ion

PO

43-

phosphate ion

• Naming ionic compounds

– The name of the cation is given first followed by

the name of the anion

• Cations:

– If the metal forms only one positive ion, the cation name is the English name for the metal

– If the metal forms more than one positive ion, the cation name is the English name followed, without a space, by the numerical value of the charge written as a Roman numeral in parentheses (this is for the Stock system)

• Anions:

– For monoatomic anions, the name is created by adding the “–ide” suffix to the stem name for the element.

• Summary of Properties

– Hardness and brittleness

• Molecular compounds tend to be soft and easily crushed because the attractions between molecules are weak and molecules can slide past each other • Ionic compounds are hard and brittle because of the

• Melting points

– To melt the particles in the solid must have

sufficient kinetic energy to overcome the

attractions between particles

• Molecular compounds tend to have weak attractions between particles and so tend to have low melting points

– Many molecular compounds are gases at room temperature

• Ionic compound tend to have strong attractions so they have high melting points

• Electrical conductivity requires the

movement of electrical charge

• Ionic compounds:

–

Do not conduct electricity in the solid state

–

Do conduct electricity in the liquid state

• The ions are free to move in the liquid state

• Molecular compounds:

–

Do not conduct electricity in the solid or liquid

state