Master Naming Chemical Compounds for DAT Exam Success

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 compounds

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

Number of oxygen atom(s)	Prefix	Suffix	Examples
One	Hypo -	- ite	BrO^- : Hypobromite
Two		- ite	BrO₂^-: Bromite
Three		- ate	BrO₃^- : Bromate
Four	Per -	- ate	BrO₄^- : Perbromate

2. Naming acids

a. Oxoacids (hydrogen + oxoanion)

Number of oxygen atom(s)	Prefix	Suffix	Examples
One	Hypo -	- ous acid	HClO : Hypochlorous acid
Two		- ous acid	HClO₂: Chlorous acid
Three		- ic acid	HClO₃ : Chloric acid
Four	Per -	- ic acid	HClO₄ : Perchloric 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	HCO₃	Hydrogen carbonate (bicarbonate)	H₂CO₃	Carbonic acid
	HSO₄	Hydrogen sulfate (bisulfate)	H₂SO₄	Sulfuric acid
	H₂PO₄	Dihydrogen phosphate	H₃PO₄	Phosphoric acid
	CH₃CO₂	Acetate	CH₃COOH	Acetic acid
	CN	Cyanide	HCN	Hydrocyanic acid
	NO₂	Nitrite	HNO₂	Nitrous acid
	NO₃	Nitrate	HNO₃	Nitric acid
	OH	Hydroxide
	ClO	Hypochlorite	HClO	Hypochlorous acid
	ClO₂	Chlorite	HClO₂	Chlorous acid
	ClO₃	Chlorate	HClO₃	Chloric acid
	ClO₄	Perchlorate	HClO₄	Perchloric acid
	MnO₄	Permanganate	HMnO₄	Permanganic acid
	BrO^-	Hypobromite	HBrO	Hypobromous acid
	BrO₂^-	Bromite	HBrO₂	Bromous acid
	BrO₃^-	Bromate	HBrO₃	Bromic acid
	BrO₄^-	Perbromate	HBrO₄	Perbromic acid
Double charged anions	HPO₄²	Hydrogen phosphate	H₃PO₄	Phosphoric acid
	CO₃²	Carbonate	H₂CO₃	Carbonic acid
	CrO₄^2-	Chromate	H₂CrO₄	Chromic acid
	Cr2O7²⁻	Dichromate	H2Cr2O7	Dichromic acid
	O₂²	Peroxide	H₂O₂	Hydrogen peroxide
	SO₃²	Sulfite	H₂SO₃	Sulfurous acid
	SO₄²	Sulfate	H₂SO₄	Sulfuric acid
	S₂O₃²	Thiosulfate	H₂S₂O₃	Thiosulfuric acid
Triple charged anions	PO₄³	Phosphate	H₃PO₄	Phosphoric acid

** NH₄⁺: Ammonium

4. Naming compounds

Fig. 1 Classification of groups

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Combinations	Naming compounds

	• Normal naming cations • Suffix “-ide” for anions • No naming for moles of cations and anions • Ex/ MgCl₂: magnesium chloride Al₂O₃: aluminum oxide Na₂O: sodium oxide **HF: hydrogen fluoride	• Normal naming cations • Normal naming polyatomic anions • No naming for moles of cations and polyatomic anions • Ex/ Mg(OH)₂: magnesium hydroxide Na₂CrO₄: sodium chromate K₂CO₃:potassium carbonate

	• Normal naming of cations • Suffix “-ide” for anions • No naming for moles of cations, but naming for moles of anions • Ex/ SnF₄: tin(IV) fluoride Fe₂O₃: iron(III) oxide	• Normal naming of cations • Normal naming polyatomic anions • No naming for moles of cations, but naming for moles of polyatomic anions • Ex/ Cu(ClO₂)₂: copper(II) chlorite Fe(ClO₄)₃: iron(III) perchlorate
	• Suffix “-ide” for anions • Use Greek prefixes to indicate the number of atoms • Omitting “mono” for first atom • Ex/ NO: nitrogen monoxide N2O4: dinitrogen tetroxide

B. Chemical Bonds

Chemical Bonds	Features
Ionic bonds	• Electrostatic attraction between oppositely charged ions in an ionic lattice. • Ionic solids are hard and brittle and generally have high melting points. • They do not conduct as solids, but conduct when molten or dissolved if soluble. • Lattice energy magnitude increases with greater ionic charge and smaller ion size. • Ex/ NaCl, MgO, CaF2
Metallic bonds	• Metallic bonding is attraction between metal cations/cores and delocalized valence electrons. • Metals generally conduct heat/electricity and are malleable and ductile. • Melting points vary widely. • Ex/ Cu, Fe, Zn, Hg
Covalent bonds (molecules)	• Covalent bonds form by sharing electron pairs, usually between nonmetals. • Bond strength: single < double < triple; bond length: single > double > triple. • Molecular substances are usually poor electrical conductors. • Intermolecular forces are distinct from covalent bonds within molecules. • Ex/ HCl, CO2, NF3

C. Covalent bond

1. Electronegativity (EN)

• Electronegativity is an atom’s ability to attract shared electrons toward itself in a covalent bond.

• Higher electronegativity generally corresponds to stronger attraction for bonding electrons and greater nonmetallic character.

• A larger electronegativity difference generally produces a more polar bond.

• 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

• Bond polarity generally increases as the electronegativity difference increases.
• A very large difference favors ionic character; bonding exists on a continuum.

Nonpolar covalent bond

• A bond between identical atoms has ΔEN = 0 and is nonpolar (H2, F2, Cl2).
• C–H bonds are commonly treated as essentially nonpolar in introductory chemistry.

• Ex/ The order of polarity between two atoms:

NaF > LiF > HF > CH > F₂

(4.0-0.9=3.1) (4.0-1.0=3.0) (4.0-2.2=2.8) (2.5-2.2=0.3) (4.0-4.0=0.0)

NaF and LiF are ionic compounds; HF has a strongly polar covalent bond, C–H is weakly polar, and F2 is nonpolar.

3. Dipole moments of polar and nonpolar molecules

Classification

Features

Polar

molecules

• Polar molecule: vector sum of bond dipoles ≠ 0.
• Ex/ H2O, NH3

Nonpolar molecules

• Nonpolar molecule: net molecular dipole = 0 because bond dipoles cancel or bonds are nonpolar.
• Ex/ CO2, CH4, BF3, BeF2

4. Molecular forces (van der Waals force)

: Attractive forces between molecules

Classification	Features
London dispersion force	• London dispersion forces arise from instantaneous/induced dipoles and occur in ALL atoms and molecules. • Strength generally increases with polarizability, electron count, and molecular surface area. • Ex/ F2 < Cl2 < Br2 < I2.
Dipole-dipole force	• Dipole–dipole attractions occur between permanent dipoles of polar molecules. • Greater molecular dipole can strengthen this interaction when other factors are comparable.
Hydrogen bond	• Hydrogen bonding occurs when H is bonded to N, O, or F and interacts with a lone pair on N, O, or F. • It is a particularly strong dipole–dipole interaction. • Ex/ HF, H2O, NH3.
Ion-dipole	• Ion–dipole attraction occurs between an ion and a polar molecule; important in solvation of ions.

** For molecules of comparable size, intermolecular-force strength often follows:

London dispersion < dipole–dipole < hydrogen bonding (when applicable)

• Compare boiling points by considering all intermolecular forces plus molecular size/polarizability; do not use one universal ranking across unrelated series.

** London dispersion becomes increasingly important as electron count and polarizability increase.

- 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

• Diatomic molecules contain two atoms.
• Common elemental diatomics: H2, N2, O2, F2, Cl2, Br2, I2.

Coordinate covalent

• A coordinate covalent bond forms when both bonding electrons are donated by the same atom.
• Ex/ NH3 + H+ → NH4+; H2O + H+ → H3O+.

Covalent network

solid

• Network covalent solids contain extended covalent bonding.
• Generally very hard with high melting points.
• Usually poor electrical conductors; graphite is an important conducting exception.
• Ex/ diamond, graphite, SiO2.

5. Octet rule

: Many main-group atoms tend toward eight valence electrons (a noble-gas-like configuration), but important exceptions exist.

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

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.

• Period 3 and heavier central atoms can form hypervalent species with more than eight electrons assigned around the central atom.

Fig. 3 atoms that are exceptions to the octet rule

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• For DAT purposes, recognize common hypervalent examples such as PCl5, SF6, and XeF4; simple d-orbital promotion is not required as an explanation.

Ex/ PCl5, SF6, XeF4

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

: Electron domains around a central atom arrange to minimize repulsions; lone pairs generally repel more strongly than bonding pairs.

Molecular shapes	Molecular geometry	Examples	Features
Linear			• 180° • AX2; 2 bonding domains, 0 lone pairs • Ex/ CO2, BeCl2, CS2
Linear			• 180° • AX2E3; 2 bonding domains, 3 lone pairs • Ex/ XeF2, I3−
Trigonal planar			• 120° • AX3; 3 bonding domains, 0 lone pairs • Ex/ BF3, AlCl3, CH2O
Bent			• Slightly <120° • AX2E; 2 bonding domains, 1 lone pair • Ex/ SO2
Bent			• About 104.5° for H2O; generally <109.5° • AX2E2 • Ex/ H2O, H2S
Tetrahedral			• 109.5° • AX4 • Ex/ CH4, SiH4, NH4+
Trigonal pyramidal			• About 107° for NH3; generally <109.5° • AX3E • Ex/ NH3, NF3, PF3
Trigonal bipyramidal			• 90°, 120°, 180° • AX5 • Ex/ PCl5, SbF5
Seesaw			• <90°, <120°, ~180° • AX4E • Ex/ SF4
T-shaped			• ~90°, 180° • AX3E2 • Ex/ ClF3
Octahedral			• 90°, 180° • AX6 • Ex/ SF6
Square pyramidal			• ~90°, 180° • AX5E • Ex/ BrF5
Square planar			• 90°, 180° • AX4E2 • Ex/ XeF4

7. Formal charges

: Formal charge is a bookkeeping charge assigned to an atom in a Lewis structure.

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• Ex/

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Atoms	Number of valence electrons	Number of bonds (bond-order sum)	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

: Hybridization describes mixing of atomic orbitals on an atom to form hybrid orbitals used in bonding.

a. Sigma (σ) bond

• A σ bond forms by head-on overlap along the internuclear axis.

• Electron density is concentrated along the internuclear axis.

• A σ bond has cylindrical symmetry about the bond axis.

• Each bonding molecular orbital can hold a maximum of 2 electrons.

b. Pi (π) bond

• A π bond forms by side-by-side overlap of parallel p orbitals.

• π electron density lies above and below the internuclear axis.

• A π bond is generally weaker than a comparable 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
sp	2 electron domains	• Linear	• C in CO₂
sp²	3 electron domains	• Bent • Trigonal planar	• S in SO₂ • C in C₂H₄
sp³	4 electron domains	• Bent • Trigonal pyramidal • Tetrahedral	• O in H₂O • N in NH₃ • C in CH₄
sp³d	5 electron domains	• Linear • Seesaw • T-Shaped • Trigonal bipyramidal	• X in XeF₂ • S in SF₄ • Cl in ClF₃ • P in PCl₅
sp³d²	6 electron domains	• Octahedral • Square Pyramidal • Square Planar	• S in SF6 • Br in BrF5 • Xe in XeF4

• Ex/ BF₃: 3 bonds and 0 lone pair Source symbol Wingdings F0E0 3+0=3 sp²types of orbital

NH₃: 3 bonds and 1 lone pair Source symbol Wingdings F0E0 3+1=4 sp³ types of orbital

9. Resonance

: Resonance is used when one Lewis structure cannot adequately represent electron distribution; the real structure is a resonance hybrid.

• Important resonance contributors delocalize electron density and can stabilize a species.

• Resonance contributors keep the same atom connectivity and total number of electrons; only electron placement changes.

• Ex/ O₃, OCN^-, NO₃^-

D. Magnetism

Magnetism	Features
Ferromagnetism	• Ferromagnetic materials can retain strong magnetization because magnetic domains align cooperatively. • Ex/ Fe, Co, Ni.
Paramagnetism	• Paramagnetic species contain one or more unpaired electrons and are weakly attracted to a magnetic field. • Ex/ O2; many transition-metal ions.
Diamagnetism	• Diamagnetic species have all electrons paired and are weakly repelled by a magnetic field. • Ex/ N2, F2, Zn2+.

E. Unit cells

: A unit cell is the smallest repeating 3-D portion of a crystal lattice that reproduces the crystal by translation.

Unit Cell	Number of atom(s) per unit cell	Features
Primitive cubic (Simple cubic)	1	• 8 corners × 1/8 = 1 atom per unit cell. • Ex/ polonium.
Body-centered cubic	2	• 8 corners × 1/8 + 1 body-center atom = 2 atoms per unit cell. • Ex/ α-iron, tungsten, sodium.
Face-centered cubic	4	• 8 corners × 1/8 + 6 faces × 1/2 = 4 atoms per unit cell. • Ex/ copper, silver, gold, nickel, lead.

III. Compound theory

Key focus of this chapter: naming chemical compounds

A. Naming compounds

B. Chemical Bonds

Molecular shapes

Examples

D. Magnetism

E. Unit cells