III. Compound theory

Key focus of this chapter: naming chemical compounds

This chapter focuses on naming chemical compounds and gives concise summaries of the important things about covalent bond which contains electronegativity, polar and nonpolar bond, dipole moments of polar and nonpolar molecules, molecular forces, octet rule, VSEPR-Valence shell, formal charges, and hybridization in more detail.

A. Naming chemical compounds

1. Naming oxoanions (atom + oxygen(s))

Number of oxygen atom(s)	Prefix	Suffix	Examples
One	Hypo -	- ite
Two		- ite
Three		- ate
Four	Per -	- ate

2. Naming acids

a. Oxoacids (hydrogen + oxoanion)

Number of oxygen atom(s)	Prefix	Suffix	Examples
One	Hypo -	- ous acid
Two		- ous acid
Three		- ic acid
Four	Per -	- ic acid

** H₂SO₃ : Sulfurous acid

H₂SO₄ : Sulfuric acid

H₃PO₄ : Phosphoric acid

b. Hydrogen + anion

Prefix	Suffix	Examples
Hydro -	- ic acid	HCl(aq): hydrochloric acid HBr(aq): hydrobromic acid HI(aq): hydroiodic acid HCN(aq): hydrocyanic acid

3. Table of naming polyatomic anions and acids

Poly atomic anions		Acids
Single charged anions	Hydrogen carbonate (bicarbonate)		Carbonic acid
	Hydrogen sulfate (bisulfate)		Sulfuric acid
	Dihydrogen phosphate		Phosphoric acid
	Acetate		Acetic acid
	Cyanide		Hydrocyanic acid
	Nitrite		Nitrous acid
	Nitrate		Nitric acid
	Hydroxide
	Hypochlorite		Hypochlorous acid
	Chlorite		Chlorous acid
	Chlorate		Chloric acid
	Perchlorate		Perchloric acid
	Permanganate		Permanganic acid
	Hypobromite		Hypobromous acid
	Bromite		Bromous acid
	Bromate		Bromic acid
	Perbromate		Perbromic acid
Double charged anions	Hydrogen phosphate		Phosphoric acid
	Carbonate		Carbonic acid
	Chromate		Chromic acid
	Dichromate		Dichromic acid
	Peroxide		Hydrogen peroxide
	Sulfite		Sulfurous acid
	Sulfate		Sulfuric acid
	Thiosulfate		Thiosulfuric acid
Triple charged anions	Phosphate		Phosphoric acid

** NH₄⁺: Ammonium

4. Naming compounds

Fig. 1 Classification of groups

Combinations	Naming compounds

B. Chemical Bonds

Chemical Bonds	Features
Ionic bonds
Metallic bonds
Covalent bonds (molecules)

C. Covalent bond

1. Electronegativity (EN)

Atomic force to pull apart the shared electrons in a covalent bond.
The atoms of higher electronegativity have strong nonmetallic character, and they easily become anions.
A greater difference of the electronegativity between two atoms means a higher polarity.
In the same period, electronegativity increases as the atomic number increases.
In the same group, electronegativity decreases as the atomic number increases.

Fig. 2 Electronegativity in the periodic table

2. Polar and nonpolar covalent bond

Classification	Features
Polar covalent bond	• The bigger difference of electronegativity between two atoms in molecules, the bigger polarity of the molecules • Big difference of electronegativity between two atoms → the higher polar molecule → ionic bonds
Nonpolar covalent bond	• The difference of electronegativity between diatomic elements is zero (H2 ,F2, Cl2) • Usually C-H bond is considered to be nonpolar.

Ex/ The order of polarity between two atoms:

Therefore, ionic bonds (NaF, LiF) have a higher polarity than covalent bonds (HF, CH, F₂)

3. Dipole moments of polar and nonpolar molecules

Classification	Features
Polar molecules
Nonpolar molecules

4. Molecular forces (van der Waals force)

: attractive forces between molecules

Classification	Features
London dispersion force
Dipole-dipole force
Hydrogen bond
Ion-dipole

** Order of strength between molecules when they have a similar molecular weight

London dispersion force < Dipole – dipole force < Hydrogen bond

** The boiling and melting point of molecules are usually dependent on London dispersion force (molecular size or weight) rather than dipole-dipole force.

– the strength order of dipole-dipole force : HI < HBr < HCl

– the strength order of London dispersion force: HCl < HBr < HI

– the strength order of the boiling and melting point: HCl < HBr < HI

** Other covalent bonds

Diatomic molecules
Coordinate covalent
Covalent network solid

5. Octet rule

: atoms in the main group tend to be positioned with 8 electrons in their outmost shell like noble gases.

a. Valence electron:

Electrons of atoms in the outermost shell
Forming covalent bonds
Main factor to distinguish the chemical properties of atoms
Ex/

Atoms	Number of valence electrons	Forming bonds
C N O F	4 5 6 7	4 3 2 1

b. Lewis structure

Repulsive and attractive forces make optimum distance between two atoms

– repulsive forces between two different atom’s electrons

– attractive forces between one atom’s electrons and another atom’s nucleus

Indicating covalent bonds

– double bonds by sharing 4 electrons (2 pairs): CO₂, O₂

– triple bonds by sharing 6 electrons (3 pairs): HCN, N₂

c. Exception to octet rule

Some atoms (Be, B, Al) have less than 8 electrons in their outermost shell.
Some atoms (past the second row in the periodic table: Si, P, Ge, Sn,…), have more than 8 electrons in their outermost shell.

Fig. 3 atoms that are exceptions to the octet rule

By using low energy in d orbitals the atoms could expand their valence shell to fill out more than 8 electrons

6. VSEPR-Valence shell (valence-shell electron-pair repulsion model)

: electrons will be positioned as far as possible from one another to maintain a stable state of molecular geometry by the repulsive forces.

Molecular shapes	Molecular geometry	Examples	Features
Linear
Linear
Trigonal planar
Bent
Bent
Tetrahedral
Trigonal pyramidal
Trigonal bipyramidal
Seesaw
T-shaped
Octahedral
Square pyramidal
Square planar

7. Formal charges

: atomic charges in a molecule.

Atoms	Number of valence electrons	Binding number	Number of non-bonding electrons
N	5	3	2
C	4	4	0
S	6	1	6

Formal charge of N: 5 −3 − 2 = 0

Formal charge of C: 4 − 4 − 0 = 0

Formal charge of S: 6 − 1 − 6 = -1

8. Hybridization

: new orbitals by overlapping atomic orbitals.

a. Sigma(σ) bond

Axis bond by head-on orbital overlap
Straight orbital overlap
Round shape by vertical section
Occupying maximum 2 electrons

b. Pi(π) bond

Sideways overlap
Dumbbell-shape by vertical section
Weaker than sigma(σ) bond

c. Application

Single bond: 1 sigma bond and 0 pi bond
Double bond: 1 sigma bond and 1 pi bond
Triple bond: 1 sigma bond and 2 pi bonds

d. Hybrid orbitals

Classification	Total number of bonds & lone pairs	Shape of molecules	Examples
		• Linear
		• Bent • Trigonal planar
		• Bent • Trigonal pyramidal • Tetrahedral
		• Linear • Seesaw • T-Shaped • Trigonal bipyramidal
		• Octahedral • Square Pyramidal • Square Planar

BF₃: 3 bonds and 0 lone pair → 3+0=3 → sp²types of orbital

NH₃: 3 bonds and 1 lone pair → 3+1=4 → sp³ types of orbital

9. Resonance

: molecule structures, which are not represented by Lewis structure, show two or more valence-bond structures by moving their electrons.

More resonance structures, greater stabilization
Maintaining same number of unpaired electrons
Ex/ O₃, OCN^–, NO₃^–

D. Magnetism

Magnetism	Features
Ferromagnetism
Paramagnetism
Diamagnetism

E. Unit cells

: the basic unit of crystal structure arranged by atoms is expressed by lattice, which is the center of an atom with 3-D repeated array.

Unit Cell	Number of atom(s) per unit cell	Features
Primitive cubic (Simple cubic)	1	• 1/8 of each atom in unit cell • Ex/ Polonium
Body-centered cubic	2	• 1 atom at center with 1/8 of each corner atom • Ex/ Tungsten, iron, sodium
Face-centered cubic	4	• 1/2 of each face-centered atoms with 1/8 of each corner atom • Ex/ Copper, silver, gold, nickel, lead, diamond