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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
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
