Chemistry B11
Chapter 3
Chemical Bonds
Chemical Bonds
1. Ionic bonds
2. Covalent bonds
3. Metallic bonds
4. Hydrogen bonds
5. Van der Waals forces
Chemical Bonds
1. Ionic bonds
2. Covalent bonds
Review
Shell 1
Maximum 2 electrons in valence shell
Hydrogen and Helium
Main-group elements
1A – 8A
Other Shells
Maximum 8 electrons in valence shell
Octet rule
Goal of atoms
Filled valence shell
Noble gases
(Stable)
+ e-
Na: 1s 2s 2p 3s
2
2
6
1
Na+: 1s 2s 2p
2
2
6
Ne: 1s 2s 2p
2
2
6
Ar: 1s 2s 2p 3s 3p
2
2
6
2
6
Octet rule
Goal of atoms
Noble gases
(Stable)
Filled valence shell
+ e-
Na: 1s 2s 2p 3s
2
2
6
Na+: 1s 2s 2p
1
2
2
Ne: 1s 2s 2p
6
2
2
6
+ e-
Cl:
1s2 2s2 2p6 3s2 3p5
-
Cl : 1s 2s 2p 3s 3p
2
2
6
2
6
Ar: 1s 2s 2p 3s 3p
2
2
6
2
6
Octet rule
Goal of atoms
Filled valence shell
Noble gases
(Stable)
+ 2e-
Mg: 1s 2s 2p 3s
2
2
6
2
Mg2+: 1s 2s 2p
2
2
6
Ne: 1s 2s 2p
2
2
6
Octet rule
Goal of atoms
Filled valence shell
Noble gases
(Stable)
+ 2e-
Mg: 1s 2s 2p 3s
2
2
6
2
Mg2+: 1s 2s 2p
2
2
Ne: 1s 2s 2p
6
2
2
6
+ 2e-
O:
1s2 2s2 2p4
2-
O : 1s 2s 2p
2
2
6
Ne: 1s 2s 2p
2
2
6
Metals: lose 1, 2 or 3 e-
Cation (Y+)
Ions
Nonmetals: gain 1, 2 or 3 e-
Anion (X-)
Number of protons and neutrons in the nucleus remains unchanged.
Cation (Y+):
Na+
Anion (X-):
Cl-
Li+
Ca2+ Al3+
F-
O2-
1A
2A
3A 4A 5A 6A 7A 8A
Transition elements
Two problems of Octet rule
1. The octet rule cannot be used for transition and inner transition elements.
Fe2+
Fe3+
Cu1+
Cu2+
2. Ions of period 1 and 2 elements with charges greater than +2 are unstable.
C
C4+
C4-
unstable
B
B3+
unstable
Naming Monatomic Cations
International Union of Pure and Applied Chemistry (IUPAC)
systematic names
Name of the metal + “ion”
H+
Li+
Hydrogen ion
Lithium ion
Cu1+
Cu2+
Copper(I) ion
Copper(II) ion
Fe2+
Fe3+
Iron(II) ion
Iron(III) ion
Hg+
Hg2+
Mercury(I) ion
Mercury(II) ion
Sn2+
Sn4+
Tin(II) ion
Tin(IV) ion
Ca2+
Al3+
Calcium ion
Aluminum ion
Naming Monatomic Cations
common name
Name of the metal +
“-ous” smaller charge
“-ic”
larger charge
Cu1+
Cu2+
Copper(I) ion
Copper(II) ion
Hg+
Hg2+
Mercury(I) ion
Mercury(II) ion
Mercurous ion
Mercuric ion
Fe2+
Fe3+
Iron(II) ion
Iron(III) ion
Ferrous ion
Ferric ion
Sn2+
Sn4+
Tin(II) ion
Tin(IV) ion
Stannous ion
Stannic ion
Cuprous ion
Cupric ion
Naming Monatomic Anions
Stem part of the name + “-ide”
Anion
Stem name
Anion name
F-
fluor
Fluoride ion
Cl-
chlor
Chloride ion
Br-
brom
Bromide ion
I-
iod
Iodide ion
O2-
ox
Oxide ion
S2-
sulf
Sulfide ion
P3-
phosph
Phosphide ion
N3-
nitr
Nitride ion
Naming Polyatomic Ions
Cation: NH4+
Ammonium
Anion:
OH-
Hydroxide
NO2-
Nitrite
NO3-
Nitrate
SO32-
Sulfite
SO42HSO3-
HSO4-
Sulfate
Hydrogen Sulfite
(bisulfite)
Hydrogen sulfate
(bisulfate)
MnO4-
Permanganate
CrO42-
Chromate
Cr2O72-
Dichromate
CO32-
Carbonate
HCO3-
Hydrogen Carbonate
(bicarbonate)
PO33-
Phosphite
PO43-
Phosphate
HPO42-
Hydrogen phosphate
H2PO4-
Dihydrogen phosphate
Ionic bonds
Metal-Nonmetal
Cl: 1s2 2s2 2p6 3s2 3p5
Na: 1s2 2s2 2p6 3s1
Na+:
-
Cl : 1s2 2s2 2p6 3s2 3p6
1s2 2s2 2p6
Cation
Anion
Sodium (Na)
NaCl
Chlorine (Cl)
matter are neutral (uncharged):
total number of positive charges = total number of negative charges
Na+ Cl-
NaCl
Ca2+ Cl-
CaCl2
Al3+ S2-
Al2S3
Ba2+ O2-
Ba2O2
Molecule of NaCl
Formula of NaCl
BaO
matter are neutral (uncharged):
total number of positive charges = total number of negative charges
Na+
NO3-
NaNO3
Ca2+
CO32-
Ca2(CO3)2
Al3+
SO42-
Al2(SO4)3
Mg2+
NO2-
Mg(NO2)2
Ca(CO3)
Naming Binary Ionic compounds
name of metal (cation) + name of anion
NaCl sodium chloride
CaO calcium oxide
Cu2O copper(I) oxide
CuO copper(II) oxide
cuprous oxide
cupric oxide
Naming Polyatomic Ionic compounds
BaCO3
barium carbonate
Li2SO4
lithium sulfate
Li2SO3
lithium sulfite
Covalent bonds
Nonmetal-Nonmetal
Metalloid-Nonmetal
Sharing of
valence electrons
Lewis Dot Structure
H
Li
He
Al
C
N
Cl
Lewis Structure
H H
Or
H
H
Cl
Or
Cl
H
H
Cl: 1s2 2s2 2p6 3s2 3p5
H: 1s1
Ar: 1s2 2s2 2p6 3s2 3p6
He: 1s2
Unshared pair of electrons
(nonbonding pair of electrons) - (Lone pair)
Cl H
x
Shared pair of electrons
(bonding pair of electrons)
Only valance electrons are involved in bonding (ionic and covalent bonds).
Electronegativity
A measure of an atom’s attraction for the electrons
Electronegativity
Ionization energy
Covalent bonds
Nonpolar covalent bond: electrons are shared equally.
Polar covalent bond: electrons are shared unequally.
δ+
δ-
H
Cl
Dipole
Electronegativity & bonds
Electronegativity Difference Between Bonded Atoms
Type of Bond
Less than 0.5
Nonpolar Covalent
0.5 to 1.9
Polar Covalent
Greater than 1.9
Ionic
H
H
2.1 – 2.1 = 0
Nonpolar covalent
N
H
3.0 – 2.1 = 0.9
polar covalent
Na
F
4.0 – 0.9 = 3.1
Ionic
Covalent compounds
CH4
H
H C H
H
H
H–C–H
H
O
CH2O
NH3
O
O
H C H H–C–H
H
H N H
O
H C H H–C–H
Correct
H
H–N–H
H H
C2H4
H
H
H C C H
H C
H
H
C–C
H
H
H
H
C=C
H
H
C H
Correct
C2H2
H C C H
H C
C H
H–C–C–H
H–C
C–H
Naming Binary Covalent compounds
Mono – Di – Tri – Tetra – Penta – Hexa – Hepta – Octa – Nona – Deca
1. Don’t use “mono” for the 1st element.
2. Drop the “a” when followed by a vowel.
prefix and full name of the first element in formula + prefix and the anion name
of the second element + “ide”
NO2
nitrogen dioxide
N2O4 dinitrogen tetroxide
CCl4 carbon tetrachloride
S2O3 disulfur trioxide
VSEPR Model
VSEPR: Valence-Shell Electron-Pair Repulsion method
Bond angle: angle between two atoms bonded to a central atom.
Each region of electron likes to be
as far away as possible from the others.
Regions of electron density
Four regions of electron density around an atom:
Bond Angles in covalent molecules
Linear molecules
2 regions
Trigonal planar
molecules
3 regions
Tetrahedral molecules
4 regions
Unshared electron paires
CH4
NH3
H2O
Polarity
1. Molecule has polar bonds.
2. Its centers of δ+ and δ- lie at different places (sides).
H δ+
δ-
δ+
δ-
O=C=O
δ+
δ+
H – Cδ-– H
H δ+
nonpolar molecule
δ-
H
O
δ-
C
N
δ+
H
polar molecule
H
δ+
H
H