Comprehensive Study Notes on Bonding, Molecular Shape, and Polarity

Elements and Physical States

• Atomic Elements (1 atom): These consist of single atoms and include metals and noble gases from Group $8A$.
• Molecular Elements: These consist of non-metals and metalloids.
  • Diatomic Elements (2 atoms): This group includes the halogens (Group $7A$), as well as $O_2$, $N_2$, and $H_2$.
  • Polyatomic Elements (3 or more atoms): Examples provided are sulfur ($S_8$) and phosphorus ($P_4$).
• Physical State Symbols:
  • $(s)$: solid
  • $(l)$: liquid
  • $(g)$: gas
  • $(aq)$: aqueous (dissolved in water)
• Elements at Room Temperature:
  • Metals: Typically solids $(s)$. Examples: $Na(s)$, $Al(s)$, $Zn(s)$.
    • Exception: Mercury ($Hg$) is a liquid: $Hg(l)$.
  • Non-metals:
    • Gases $(g)$: Noble gases (Group $8A$), $H_2$, $N_2$, $O_2$, $F_2$, and $Cl_2$.
    • Liquids $(l)$: Bromine ($Br_2$) is a liquid: $Br_2(l)$.
    • Solids $(s)$: Most other non-metals and metalloids, such as iodine ($I_2(s)$).
• Allotropes: Multiple distinct chemical formulas of an element existing in the same physical state.

Ionic and Covalent Compounds

• Compound: A substance containing two or more different elements.
• Ionic Compounds:
  • Comprised of a positively charged cation ($+$) and a negatively charged anion ($-$).
  • Typically composed of a metal cation ($+$) and a non-metal anion ($-$).
  • Ionic Bond: A stabilizing electrostatic interaction between a cation and an anion. These bonds are very strong.
  • Formula Unit: The smallest repeating set of ions in an ionic compound.
• Polyatomic Ion: A charged molecule containing two or more atoms that are covalently bonded.
• Covalent (Molecular) Compound:
  • Two or more atoms of different elements covalently bonded together.
  • Composed of two or more non-metals.
  • These compounds are neutral and contain no charged ions.
  • Covalent Bond: The sharing of outermost electrons between two atoms. These bonds are stable but generally weaker than ionic bonds. Shared electrons are stabilized by the nucleus of each participating atom.

Electronegativity and Chemical Bonding

• Electronegativity ($EN$): A measurement of how readily a nucleus attracts electron density within a chemical bond.
  • Electronegative elements (non-metals) prefer to form anions.
  • Less electronegative elements (metals) prefer to form cations.
• $\Delta EN$ (Electronegativity Difference): The difference in electronegativity between two atoms determine the bond type:
  • $\Delta EN < 0.4$: Pure covalent (nonpolar) bond.
  • $0.4 < \Delta EN < 2.0$: Polar covalent bond.
  • $2.0 < \Delta EN$: Ionic bond.

Periodic Trends and Ion Formation

• Cation Formation: Metals typically lose electrons to generate positive cations ($+$).
  • Main group metals (A) form cations with a $\text{charge} = \text{group number}$.
• Anion Formation: Non-metals typically gain electrons to generate negative anions ($-$).
  • Non-metals form anions with a $\text{charge} = \text{group number} - 8$.
• Transition Metals (B): Typically adopt a $+2$ charge but can adopt a range of positive charges.
  • Fixed Charge Transition Metals (To Memorize): Silver ($Ag$), Cadmium ($Cd$), and Zinc ($Zn$) exclusively form $Ag^+$, $Cd^{2+}$, and $Zn^{2+}$.
• Predicting Ion Charges (Exercises):
  • $Mg$: $+2$
  • $S$: $-2$
  • $Ag$: $+1$
  • $N$: $-3$
  • $Br$: $-1$
  • $Li$: $+1$
  • $Zn$: $+2$
  • $Sc$: $+3$

Valence Electrons and the Octet Rule

• Valence Electrons: Electrons in the highest energy principle energy level (largest $n$); these are the most readily accessible and highest energy electrons.
• Octet Rule: Atoms gain, lose, or share electrons to reach noble gas configurations, which consist of $8$ valence electrons.
• Main Group Elements: The number of valence electrons is equal to the group number.
• Abbreviated Electron Configurations: Used to identify valence and core electrons.
  • $Cl$: $[Ne] 3s^2 3p^5$ ($7$ valence electrons)
  • $Br$: $[Ar] 4s^2 3d^{10} 4p^5$ ($7$ valence electrons)
  • $Li$: $[He] 2s^1$ ($1$ valence electron)
  • $C$: $[He] 2s^2 2p^2$ ($4$ valence electrons)
• Lewis Dot Structures: A method to abbreviate valence electrons using dots ($1$ dot = $1$ valence electron).
  • Elements in the same group have the same number of valence dots.
• Ionization and Noble Gas Configurations:
  • $1A-3A$ elements form cations ($Na \rightarrow Na^+$).
  • $5A-7A$ elements form anions ($Cl \rightarrow Cl^-$).
  • Cations and anions achieve a noble gas configuration.
  • Transition Metal Ionization: Typically lose the outermost highest energy $s$ orbital electrons first.
    • Example: $Co [Ar] 4s^2 3d^7 \rightarrow Co^{2+} [Ar] 3d^7$.
    • In many cases, loss of both outermost $s$ and $d$ orbital electrons occurs.

Writing Lewis Structures and Configurations for Ions

• Cations: Subtract electrons equal to the positive charge.
  • $K = [Ar] 4s^1 \text{ (19 electrons)} \rightarrow K^+ = [Ar] \text{ (18 electrons)}$.
• Anions: Add electrons equal to the negative charge.
  • $Cl = [Ar] 3s^2 3p^5 \text{ (17 electrons)} \rightarrow Cl^- = [Ar] 3s^2 3p^6 \text{ (18 electrons)}$.
• Covalent Bond Details:
  • Shared electrons are counted as belonging to both atoms.
  • Lone pair (nonbonding) electrons: Not shared; belong only to the atom they are on.
  • Example $H_2O$: Oxygen has $4$ shared electrons and $4$ lone pair electrons, totaling $8$. Hydrogen has $2$ shared electrons ($duet$).
• Bond Representations: Written as lines.
  • $1$ line = $1$ bond = $2$ shared electrons (Single bond, e.g., $F-F$).
  • $2$ pairs shared = Double bond (e.g., $O=O$).
  • $3$ pairs shared = Triple bond (e.g., $N \text{ (triple bond) } N$).
  • Exception: Hydrogen forms a duet ($2$ electrons) to reach the configuration of $He$ ($1s^2$).

Detailed Steps for Drawing Lewis Structures

1. Count total valence electrons ($ve^-$) for all atoms.
2. If cation: subtract electrons equal to charge.
3. If anion: add electrons equal to charge.
4. Determine the central atom: Hydrogen is never central. The central atom is the least electronegative atom.
5. Arrange atoms symmetrically around the central atom.
6. For $n=1, n=2$ elements: place a maximum of $4$ atoms around each central atom.
7. Form a single bond ($2$ electrons, $1$ line) between central and surrounding atoms.
8. Subtract $2$ electrons from the total for each bond made.
9. Place electron pairs ($2$ dots) around outer atoms until octets are satisfied (prioritize most electronegative first).
10. If an octet is unfulfilled: convert lone pairs into double/triple bonds only if needed.
11. Verify all atoms have complete octets.

Examples of Lewis Structure Calculations:

• $CBr_2I_2$: $4 + (2 \times 7) + (2 \times 7) = 32 ve^-$
• $CO_3^{2-}$: $4 + (3 \times 6) + 2 = 24 ve^-$. (Requires brackets and charge indicator).
• $CCl_4$: $4 + (4 \times 7) = 32 ve^-$
• $HCN$: $1 + 4 + 5 = 10 ve^-$
• $SO_3$: $6 + (3 \times 6) = 24 ve^-$
• $SeF_2$: $6 + (2 \times 7) = 20 ve^-$
• $CH_2O$: $4 + (2 \times 1) + 6 = 12 ve^-$
• $SO_2$: $6 + (2 \times 6) = 18 ve^-$
• $H_3O^+$: $(3 \times 1) + 6 - 1 = 8 ve^-$
• $CN^-$: $4 + 5 + 1 = 10 ve^-$. Structure: $[C \text{ (triple bond) } N]^-$

Formal Charges and Stability

• Formula: FC_{atom} = (\text{Group Number of atom}) - (\text{# of bonds} + \text{# of lone pair electrons}).
• Favorable Structures:
  1. Small formal charge values ($0$ is best).
  2. Fewest atoms with non-zero formal charges.
  3. Negative formal charges on more electronegative atoms; positive on less electronegative.
  4. No adjacent like charges (positive-positive or negative-negative).
• The total charge of the molecule must equal the sum of all formal charges.

Lewis Structures by Analogy (Isoelectronic Species)

• Isoelectronic: Atoms or ions with the same number of valence electrons can be exchanged to create valid structures.
• Common Isoelectronic Sets:
  • $4 ve^-$: $C$, $Si$, $N^+$, $O^{2+}$.
  • $5 ve^-$: $N$, $P$, $O^+$, $C^-$.
  • $6 ve^-$: $O$, $S$, $Cl^+$.
  • $7 ve^-$: $F$, $Cl$, $O^-$, $S^-$, $N^{2-}$.
• Examples by Analogy:
  • $CH_4$ and $NH_4^+$
  • $O_3$ and $SO_2$
  • $CO_3^{2-}$ and $NO_3^-$
  • $CO_2$ and $NO_2^+$

VSEPR Theory and Molecular Geometry

• VSEPR (Valence Shell Electron Pair Repulsion) states that electron pairs (bonds and lone pairs) repel each other and arrange to minimize this repulsion.
• Charge Sphere/Electron Group: Either a single lone pair or a bonded atom (single, double, and triple bonds count as $1$ group).
• Steric Number ($SN$): SN = (\text{# bonded atoms}) + (\text{# lone pairs}).
• Electron Pair Geometry: Based on total charge spheres.
  • $SN = 2$ (Linear): Bond angle = $180^{\circ}$.
  • $SN = 3$ (Trigonal Planar): Bond angle = $120^{\circ}$.
  • $SN = 4$ (Tetrahedral): Bond angle = $109.5^{\circ}$.
• Molecular Geometry: Accounts for the specific repulsion of lone pairs. Lone pairs take up more space, distorting and shrinking bond angles.
  • $4$ spheres, $0$ lone pairs: Tetrahedral ($109.5^{\circ}$).
  • $4$ spheres, $1$ lone pair: Trigonal Pyramidal ($107.5^{\circ}$).
  • $4$ spheres, $2$ lone pairs: Bent ($105.5^{\circ}$).
  • $3$ spheres, $0$ lone pairs: Trigonal Planar ($120^{\circ}$).
  • $3$ spheres, $1$ lone pair: Bent ($116.0^{\circ}$).
• Practice Examples:
  • $CH_4$: Tetrahedral molecular geometry.
  • $H_2Se$: Bent molecular geometry (based on Tetrahedral electron geometry).
  • $PH_3$: Trigonal Pyramidal molecular geometry.

Molecular Polarity

• Net Polarity: Determined by individual bond dipoles and molecular shape.
• Vector Analysis: Dipoles are vectors (magnitude and direction). If they don't cancel, the molecule has a net dipole and is polar.
• Symmetry and Cancellation: Dipoles cancel only if:
  1. Bonds are symmetrically arranged.
  2. Bonded atoms are identical (same electronegativity).
• Polarity Shortcut: A molecule is nonpolar if:
  1. All atoms bonded to the center are identical OR $\Delta EN < 0.4$.
  2. The central atom has no lone pairs and has a symmetrical geometry (linear, trigonal planar, tetrahedral, square planar).
  • Otherwise, the molecule is polar.
• Examples:
  • $CO_2$: Nonpolar (dipoles cancel linearly).
  • $H_2O$: Polar (Bent geometry prevents dipole cancellation).
  • $CCl_4$: Nonpolar.
  • $NH_3$: Polar.

Chemical Formulas and Nomenclature of Ionic Compounds

• Chemical Formula: Specifies the number of atoms of each element.
  • Subscript: Indicates number of atoms or polyatomic subunits (in parentheses).
  • Example: $Ca_3(PO_4)_2$ has $3$ Calcium, $2$ Phosphorus, and $8$ Oxygen atoms.
• Polyatomic Ion Naming (Suffix Rules):
  • -ate: Base ion (most common number of oxygens).
  • per- / -ate: One more oxygen than -ate ion.
  • -ite: One less oxygen than -ate ion.
  • hypo- / -ite: Two less oxygens than -ate ion.
• Predicting Ionic Formulas:
  1. Total charge of a neutral compound must be $0$.
  2. Formula: \text{charge cation} \times \text{# cation} + \text{charge anion} \times \text{# anion} = 0.
  3. Cross-over method: Place the charge of one ion as the subscript of the other.
  • Examples: $Fe^{3+}$ and $S^{2-}$ form $Fe_2S_3$. $Al^{3+}$ and $N^{3-}$ form $AlN$ (simplify subscripts).
• Naming Binary Ionic Compounds (Class A - Main Group Metals):
  1. Name of metal cation.
  2. Name of non-metal anion with -ide ending.
  • Stems: Ox- (Oxygen), Nitr- (Nitrogen), Phosph- (Phosphorus), Sulf- (Sulfur).
• Variable Charge Transition Metals (Class B):
  1. Use Roman numerals in parentheses to indicate the metal's charge.
  2. Examples: $FeSO_3$ is Iron(II) sulfite. $Cr(HCO_3)_3$ is Chromium(III) hydrogen carbonate.
  3. Fixed Charge Exceptions: Do not use Roman numerals for $Ag^+$, $Cd^{2+}$, or $Zn^{2+}$.

Nomenclature of Acids, Hydrates, and Covalent Compounds

• Hydrates: Ionic compounds with associated water molecules.
  • Naming: [Salt Name] + [Prefix]-hydrate.
  • Example: $CoCl_2 \cdot 6H_2O$ is Cobalt(II) chloride hexahydrate.
  • Anhydrous: A hydrate that has lost its water.
• Binary Acids (H + non-metal):
  • Gas Phase ($g$): Hydrogen [Stem]-ide. (e.g., $HCl(g)$ = Hydrogen chloride).
  • Aqueous Phase ($aq$): hydro-[Stem]-ic acid. (e.g., $HCl(aq)$ = Hydrochloric acid).
• Polyatomic (Oxy) Acids:
  • If polyatomic ion ends in -ite, acid ends in -ous acid.
  • If polyatomic ion ends in -ate, acid ends in -ic acid.
  • Special cases: $H_2SO_4$ (sulfuric), $H_3PO_4$ (phosphoric).
• Binary Covalent Compounds (2 non-metals):
  • Use prefixes for both elements (mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-).
  • Second element ends in -ide.
  • Example: $P_2O_5$ is Diphosphorus pentoxide. $N_2O$ is Dinitrogen monoxide.

Organic Nomenclature

• Alkanes (Saturated Hydrocarbons): Formula $C_nH_{2n+2}$. Suffix -ane.
  • $1$ Carbon: Methane
  • $2$ Carbons: Ethane
  • $3$ Carbons: Propane
  • $4$ Carbons: Butane
  • $7$ Carbons: Heptane ($C_7H_{16}$)
• Alkenes (Unsaturated Hydrocarbons): Contain a double bond. Formula $C_nH_{2n}$. Suffix -ene.
  • Example: $C_2H_4$ (Ethene), $C_3H_6$ (Propene), $C_4H_8$ (Butene).
• Alcohols: Replace one $H$ in an alkane with an $-OH$ group. Suffix -ol.
  • $CH_3OH$: Methanol
  • $CH_3CH_2OH$: Ethanol
  • $CH_3CH_2CH_2CH_2OH$: Butanol