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Chapter 4 Compounds

1. Ionic Compounds

2. Molecular Compounds (Covalent Compounds)

3. Names of Ions

4. Naming Compounds

Table of Contents

Chapter 4 Compounds

• Predict how atoms combine and make compounds.

What plays an important role in combination of atoms?

Warm up

• Predict changes during formation of compounds.

• List down some items that you use in life and classify

them in a way with same characteristics.

Chapter 4 Compounds

• A compound is a pure substance containing two or

more types of atoms in definite proportion.

• And its smallest unit is called molecule.

• Compounds do not resemble their constituents.

Chapter 4 Compounds

Compounds

Ionic Compounds Covalent Compounds

Polar Nonpolar

• According to their bond structures;

Chapter 4 Compounds

Chapter 4 Compounds

Chapter 4 1. Ionic Compounds• Ionic compounds are formed from metals and

nonmetals by transferring their valance electrons.

• Ionic compounds have ionic bonds.

• They do not conduct electricity in their solid state. Their

aqueous solutions conduct electricity.

• They are solids at room conditions.

• They have crystalline structures.

• They generally dissolve in water and produce ions.

Chapter 4 1. Ionic Compounds

Chapter 4 1. Ionic Compounds

Chapter 4 1. Ionic Compounds

Chapter 4 1. Ionic Compounds

Example 1Table salt, Magnesium fluoride, Calcium oxide, Silver

iodide, Aluminum oxide…etc.

Na Na+ + 1e-

Cl + 1e- Cl-

Na+ Cl-+ NaCltable salt

Chapter 4 2. Covalent Compounds• Covalent compounds are formed between nonmetals

by sharing their valance electrons.• Covalent compounds have covalent bonds.

Carbon dioxide.Water

Chapter 4 2. Covalent Compounds• Covalent bonds found between molecules composed of

the same atoms are nonpolar covalent bonds.

Bromine, Br2 Nitrogen gas, N2 Oxygen gas, O2

• Covalent bonds between atoms with different nonmetals are called polar covalent bonds.

HCl HF

Chapter 4 2. Covalent Compounds

• They are composed of nonmetal elements.

• They generally do not conduct electricity, because their

molecular structures are conserved while dissolving.

• They can be solid, liquid and gaseous at room

conditions.

Example 2Water, alcohol, sugar, acetic acid, ozone …etc.

Chapter 4 2. Covalent Compounds

Chapter 4 3. Names of Ions

• In order to write formula of compounds names of ions should be known.

Monoatomic Cations (Metal ions)

+1 +2H+ Hydrogen Mg+2 Magnesium

Na+ Sodium Hg+2 Mercury (II)

K+ Potassium Ca+2 Calcium

Hg+ Mercury Cu+2 Copper (II)

Ag+ Silver Ba+2 BariumCu+ Copper Ni+2 Nickel

Li+ Lithium Zn+2 Zinc

Chapter 4 3. Names of IonsMonoatomic Cations (Metal ions)

+2Fe+2 Iron (II)

Cr+2 Chromium (II)

+3Fe+3 Iron (III)

Cr+3 Chromium (III)

+4Pb+4 Lead (IV)

Sn+4 Tin (IV)

Pb+2 Lead (II)

Sn+2 Tin (II)

+1NH4

+1 Ammonium

H3O+1 Hydronium

Polyatomic Cations

Chapter 4 3. Names of IonsMonoatomic Anions (Nonmetal ions)

-1 -2F- Fluoride O-2 Oxide

Cl- Chloride S-2 Sulfide

Br- Bromide -3

I- Iodide N-3 Nitride

H- Hydride P-3 Phosphide

Chapter 4 3. Names of IonsPolyatomic Anions

-1 -2

OH-1 Hydroxide SO4-2 Sulfate

NO3-1 Nitrate SO3

-2 Sulfite

NO2-1 Nitrite CO3

-2 Carbonate

CH3COO-1 Acetate CrO4-2 Chromate

ClO2-1 Chlorite Cr2O7

-2 Dichromate

ClO3-1 Chlorate MnO4

-2 Manganate

CN-1 Cyanide C2O4-2 Oxalate

MnO4-1 Permanganate

-3PO4

-3 Phosphate PO3-3 Phosphite

Chapter 4 3. Names of IonsMonoatomic Cations (Metal ions)

Chapter 4 3. Names of IonsPolyatomic Ions

Chapter 4 3. Names of IonsPolyatomic Ions

Chapter 4 3. Names of IonsFormation of Monoatomic Ions

Chapter 4 3. Names of Ions

• A formula is a combination of symbols and numbers that represents compounds.

H3PO4

symbols of elements

numbers of atoms

• 3 different elements, H, P and O.• Contains 3-H, 1-P, 4-O atoms.• 1 molecule of H3PO4 contains a

total of 8 atoms.

• Subscript 1 is not written in the formulas of compounds.

Chapter 4 3. Names of Ions

Writing Formulas of Ionic Compounds

• Net charge of ions of a compound must be zero.

X+n

Y-m

X nYm

Example 3Write the formula of compounds between ions given below.a. K+ and Br- b. Mg+2 and O-2 c. Ca+2 and N-3

Chapter 4 3. Names of Ions

Writing Formulas of Ionic Compounds

Solutiona. KBr b. MgO c. Ca3N2

Example 4Write the formula of compounds between ions given below.a. Li+ and CO3

-2 b. Ba+2 and NO3-1 c. NH4

+ and P-3

Solutiona. Li2CO3 b. Ba(NO3)2 c. (NH4)3P

Chapter 4 3. Names of IonsFinding the Oxidation Number of an Element in a Compound

• Sum of the oxidation numbers of elements in a compound is zero.

• Oxidation numbers of some common ions like Na+, K+, Li+, Ca+2, Ba+2, Zn+2, Ag+, Al+3 are constant.

• In general oxygen has -2 oxidation number and hydrogen has +1 oxidation number.

Example 5Find the oxidation number (valency) of C in Li2CO3.

Solution

Li2CO3+1 x -2 2 Li + C + 3 O = 0

2 x (+1) + x + 3 x (-2) = 0x = +4

Chapter 4 3. Names of IonsFinding the Oxidation Number of an Element in a Compound

Chapter 4 3. Names of IonsFinding the Oxidation Number of an Element in a Compound

2 Al + 3.(S+4.O) = 02 . (+3) + 3.{x+4.(-2)} = 06 + 3.{x-8} = 06 + 3x - 24 = 03x = 18x = +6

Solution

Al2(SO4)3+3 x -2

Example 5Find the oxidation number (valency) of S in Al2(SO4)3.

Chapter 4 3. Names of IonsFinding the Oxidation Number of Elements in Polyatomic Ions

Chapter 4 4. Naming CompoundsA. Naming Ionic Compounds

Name of Metal + Name of Nonmetal ion

Example 6Name the following ionic compounds.a. NaBr b. Al2O3 c. ZnF2 d. Ba3N2

Solutiona. Sodium Bromide b. Aluminum Oxidec. Zinc Fluoride d. Barium Nitride

Chapter 4 4. Naming CompoundsA. Naming Ionic Compounds

Chapter 4 4. Naming CompoundsA. Naming Ionic Compounds

Example 7Name the following ionic compounds.a. KOH b. NiSO4 c. Zn(MnO4)2 d. NH4Cl

Solutiona. Potassium Hydroxide b. Nickel Sulfatec. Zinc Permanganate d. Ammonium Chloride

Chapter 4 4. Naming CompoundsA. Naming Ionic Compounds

Example 8Name the following ionic compounds.a. Fe2O3 b. CuSO4 c. Pb(NO3)4 d. SnC2O4

Solutiona. Iron (III) Oxide b. Copper (II) Sulfatec. Lead (IV) Nitrate d. Tin (II) oxalate

Chapter 4 4. Naming CompoundsB. Naming Molecular Compounds

Number + Name of Nonmetal + Number + Name of Nonmetal

• Greek numbers are used to show the number of atoms.

Mono 1 Hexa 6Di 2 Hepta 7Tri 3 Octa 8Tetra 4 Nona 9Penta 5 Deca 10

Cl2O5Dichloro Pentoxide

Chapter 4 4. Naming CompoundsB. Naming Molecular Compounds, continued

Chapter 4 4. Naming Compounds

Example 9Name the following molecular compounds.a. NO2 b. CCl4 c. N2O5 d. P2O3

Solutiona. Nitrogen dioxide b. Carbon tetrachloridec. Dinitrogen pentoxide d. Diphosphorus trioxide

B. Naming Molecular Compounds

Chapter 4 4. Naming Compounds

Chapter 4 4. Naming CompoundsThe Law of Definite Proportion (Proust’s Law)

• Proust stated that elements of a compound are combined in definite proportion by mass.

Example 10Find the definite proportions of elements in the following compounds.a. CH4 b. SO2 c. KBr

16 amu 12 amu 16 amu

CO2:mCmO

=1232 =

38

Chapter 4 4. Naming CompoundsThe Law of Definite Proportion (Proust’s Law)

Solution

CH4:mCmH

=124 =

31

a.

= 3

SO2 :mSmO

=3232 =

11

b.= 1

KBr :mKmBr

=3980

c.

End of the chapter 4

Chapter 2 Bonds in Solids and Liquids

1. Metallic Bonds

2. Ionic Solids

3. Network Solids

4. Dipole-Dipole Forces

5. Van der Waals Forces

6. Hydrogen Bond

Table of Contents

Chapter 2

Warm up

• List some substances in different states of matter and try to

explain why they are solid, liquid or gas.

• Remember polarity of molecules, and give some

examples.

• Which substance is known as the hardest and how it is

used?

Bonds in Solids and Liquids

Chapter 2

• In nature substances usually are found in three states.

Intermolecular forces of attractions play important role in

solids and liquids.

Bonds in Solids and Liquids

1. Metallic BondsIn metal atoms valence electrons move freely from the empty

orbitals of one atom to another. These electrons that can move

freely around the nuclei of the atoms form an “electron sea”. An

attraction force occurs between the negatively charged “sea of

electrons” and the positively charged nuclei. Metal atoms are

held together because of this attractive force. This is called the

metallic bond.

Chapter 2 1. Metallic Bonds

1s2 2s2 2p6

11Na:3s1 3p0

• One valance electron in 3s orbital freely move in 3p orbital of

another atoms.

Chapter 2 1. Metallic Bonds

Chapter 2 1. Metallic Bonds

Chapter 2 1. Metallic Bonds

Chapter 2

• In a group, metallic bond strength generally decreases

from up to down.

• In a period, strength generally increases from left to right.

• Metals are good conductors of heat and electricity.

• Metals can be drawn into wires and hammered into shape

easily.

Example 1

Compare the metallic bonds in Na, Mg and Al and explain.

1. Metallic Bonds

Chapter 2 2. Ionic Solids

• When metal and nonmetal atoms come together, they

form ionic bonds.

• Electrostatic attraction occurs between the positive and

negative charges holding the ions together.

• Metal ions are surrounded by nonmetal ions and

nonmetal ions surrounded by metal ions.

• The melting and boiling points of ionic solids are very high.

• In molten state and in solutions they conduct electricity.

Chapter 2 2. Ionic Solids

Chapter 2 2. Ionic Solids

Chapter 2 2. Ionic Solids

Chapter 2 3. Network Solids• Network solids are giant arrangements of matter in which

atoms are covalently bonded together in a continuous two or

three dimensional array. You can think of network solids as

giant molecules.

• Graphite, diamond, SiC and SiO2 are some examples.

Chapter 2 3. Network Solids

• Each carbon atom is covalently bonded to four others with

sp3 hybrid orbitals forming a tetrahedral shape.

• It is the hardest substance known.

•It does not have free electrons. Thus it cannot conduct

electricity.

Diamond

Graphite •Carbon atoms are bonded to three others with sp2 hybrid

orbitals forming hexagonal shapes with 120o angle.

• It is soft substance known.

•It can conduct electricity.

Chapter 2 3. Network Solids

Chapter 2 3. Network Solids

Chapter 2 4. Dipole-Dipole Forces

• In polar covalent substances, there is an attraction

between the positive end of one dipole and the negative end

of neighboring dipoles. This attraction is called dipole-dipole

attraction.

Chapter 2 4. Dipole-Dipole Forces

Chapter 2 5. Van Der Waals Forces• In noble gases and non polar molecules movement of

electrons causes in non polar molecules becoming

temporarily polar and an instantaneous dipole is formed.

Momentarily dipole molecule causes neighboring molecules

to become polar. Thus a weak attraction occurs between the

molecules. This attraction is called Van der Waals Forces.

• It depends upon the electron density of the atoms. It is

stronger between molecules with high molecular masses.

Example 2

Compare the boiling points of CH4, H2, N2, O2 gases.

Chapter 2 5. Van Der Waals Forces

Chapter 2 6. Hydrogen Bonds

• F, O, and N are the most electronegative elements.

Therefore their compounds with hydrogen (HF, H2O and

NH3) are highly polar. This causes an attraction force

stronger than usual dipole-dipole forces. This stronger

intermolecular forces are called hydrogen bonds.

• Unexpected increase in the boiling points of HF, H2O and

NH3 can be explained by hydrogen bond.

Chapter 2 6. Hydrogen Bonds

Chapter 2 6. Hydrogen Bonds

End of the chapter 2