Inorganic and Organic Chemistry: Chemical Bonding, Nomenclature, and Oxidation Rules for Formulas, and Oxidation States

Core Principles of Chemical Bonding

Chemical compounds are formed when two or more atoms are tied together by a force of attraction known as a chemical bond. These bonds act as the "glue" that holds particles together in matter and are responsible for giving molecules their specific properties.

Valence Electrons and the Octet Rule

• Valence Electrons: These are the electrons located in the last energy level (outermost shell) of an atom. They are generally the only electrons involved in the formation of chemical bonds. The number of valence electrons for an element can be determined by its group number on the periodic table. For example, an atom with the configuration $2-8-8-1$ has one valence electron.

• The Octet Rule: Proposed by Gilbert N. Lewis, this rule states that atoms are most stable when they have a complete set of $8$ valence electrons in their outermost shell. Elements will borrow, lose, or share electrons with other elements to achieve this stability. A helpful mnemonic is: "$2$ will do but $8$ is great!"

  • The total number of bonds an atom can make is equal to the number of electrons it needs to complete its valence shell.

  • Specific Examples:

    • Hydrogen ($H$): Atomic number is $1$. It has one electron in its first shell (which can hold a maximum of 2).

    • Carbon ($C$): Atomic number is $6$. It has two shells; the outermost shell requires more electrons to reach an octet.

    • Nitrogen ($N$): Atomic number is $7$.

    • Oxygen ($O$): Atomic number is $8$.

Lewis Dot Structures

Lewis Dot Structures are a notation showing the valence electrons surrounding an atomic symbol. These structures include paired electrons and unpaired electrons.

Rules for Atomic Lewis Dots

1. No side of the atomic symbol can have more than two dots.

2. When filling the four sides of the element symbol, each side must receive one dot before any side is doubled up.

3. Exceptions: Hydrogen and Helium do not follow the standard octet behavior in the same way as larger atoms.

Rules for Drawing Lewis Structures of Compounds

1. Count Valence Electrons: Calculate the total number of valence electrons available from all atoms in the compound.

   • Example ($CH_4$): Carbon provides $4$ valence electrons. Each of the four Hydrogen atoms provides $1$ valence electron, totaling $4$. The total count is $8$ valence electrons.

2. Identify the Central Atom: This is typically the first atom written in the formula. If the first atom is Hydrogen, the second atom becomes the central atom (Hydrogen can never be the central atom and never has more than one bond).

3. Place Terminal Atoms: Arrange terminal atoms around the central atom.

4. Complete Octets: Complete the octets for all atoms (except Hydrogen) using lone pairs of electrons.

5. Investment Principle: Each atom contributes one electron per bond but counts the bond as two electrons, essentially doubling its "investment."

6. Verification: Check the structure by counting total electrons used; they must match the initial count from Step 1. Every bond "feels" like two electrons to every element it borders.

Major Types of Chemical Bonds

1. Ionic Bonds

• Definition: An ionic bond occurs when an electron leaves one atom and enters the orbit of another exothermically. This creates two oppositely charged ions that attract one another via electrostatic forces.

• Formation: Generally formed between metals and nonmetals.

• Electronegativity: Occurs when the electronegativity difference between bonding atoms is $1.7$ or higher.

• Case Study (Rock Salt): Bonding between Sodium ($Na$) and Chlorine ($Cl$). Sodium is a silvery metal with $1$ valence electron. Chlorine is a yellow-green gas needing $1$ electron. When combined, they burn as electrons are exchanged, the metal dissolves, and the gas disappears to form a crystal with a rock salt structure.

2. Covalent Bonds

• Definition: A type of chemical bond involving the mutual sharing of electrons between two atoms.

• Classification by Bonds:

  • Single Bond: Mutual sharing of one pair of electrons.

  • Double Bond: Sharing of two pairs (e.g., $O_2$).

  • Triple Bond: Sharing of three pairs (e.g., $N_2$).

• Classification by Polarity:

  • Non-polar Bond: Shared electrons are equidistant between atoms. The electron cloud is evenly distributed. This occurs when the electronegativity difference is between $0$ and $0.5$. Example: $H_2$ ($H:H$).

  • Polar Bond: Shared electrons are closer to the more electronegative atom, resulting in a denser electron cloud in that vicinity. This occurs when the electronegativity difference is between $0.6$ and $1.6$. Example: $HCl$ ($H-Cl$).

3. Hydrogen Bonds

• Definition: The attractive force between a hydrogen atom (attached to an electronegative atom) of one molecule and an electronegative atom of a different molecule.

• Key Atoms: Usually involves Oxygen ($O$), Nitrogen ($N$), or Fluorine ($F$), which carry partial negative charges.

• Water Example: In water, Oxygen is more electronegative than Hydrogen. Oxygen has two unshared pairs of electrons and two shared pairs (covalently bonded to Hydrogen). This creates partial positive ($\\delta+$) and partial negative ($\\delta-$) charges that facilitate hydrogen bonding between molecules.

4. Metallic Bonds

• Definition: Electrostatic forces of attraction between metal cations (positively charged nuclei and inner energy level electrons) and delocalized outer energy level electrons.

• Structure: These bonds hold particles together in metals through the attraction of metal cations to the sea of shared, delocalized electrons.

Oxidation Numbers and Redox Principles

Fundamentals of Redox

• Oxidation: Loss of Electrons.

• Reduction: Gain of Electrons.

• Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain).

• Oxidising Agents: Substances that cause oxidation (e.g., $O_2$, $Cl_2$, $MnO_4^{1-}$, $Cr_2O_7^{2-}$).

• Reducing Agents: Substances that cause reduction (e.g., $C$, $CO$, $H_2$).

Rules for Assigning Oxidation Numbers (O.N.)

1. Free Elements: The O.N. is always $0$ (e.g., $Na = 0$, $P = 0$, $S_8 = 0$, $Cu = 0$, $H_2 = 0$, $Cl_2 = 0$).

2. Molecules: The sum of all oxidation numbers in a neutral molecule is $0$ (e.g., $CaCO_3$, $(NH_4)_2SO_4$).

3. Simple Ions: The O.N. equals the charge of the ion (e.g., $Cl^- = -1$, $O^{2-} = -2$, $Al^{3+} = +3$, $Na^+ = +1$, $N^{3-} = -3$).

4. Complex Ions: The sum of O.N.s equals the charge on the ion (e.g., $SO_4^{2-} = -2$, $PO_4^{3-} = -3$, $NH_4^+ = +1$, $NO_3^- = -1$).

5. Hydrogen ($H$): Usually $+1$ (e.g., $H_2O$, $HCl$, $CH_3COOH$). Exception: In metal hydrides, it is $-1$ (e.g., $NaH$, $CaH_2$).

6. Oxygen ($O$): Usually $-2$. Exceptions: In peroxides ($H_2O_2$, $Na_2O_2$, $BaO_2$), it is $-1$. When bonded to Fluorine ($F$), it is $+2$.

7. Alkali Metals (Group I): Always $+1$ in compounds ($Li$, $Na$, $K$, $Rb$, $Cs$, $Fr$).

8. Alkaline Earth Metals (Group II): Always $+2$ in compounds ($Be$, $Mg$, $Ca$, $Sr$, $Ba$, $Ra$).

9. Halogens (Group VII): Usually $-1$ ($F$, $Cl$, $Br$, $I$). Exception: When bonded to a more electronegative element (e.g., in $Cl_2O$, $Cl$ is $+1$; in $ClO_2$, $Cl$ is $+4$).

10. Direction of Change: Oxidation involves a decrease in oxidation number (Wait: Slide notes also state $Mg$ $0$ to $+2$ is oxidation; standard definition is an increase in number. Slide text explicitly states: "Oxidation is a decrease in oxidation number; Reduction is a gain of oxidation number").

    • Example: $Mg + Cu^{2+} \rightarrow Mg^{2+} + Cu$. $Mg$ changes from $0$ to $+2$ (loss of 2 electrons); $Cu$ changes from $+2$ to $0$ (gain of 2 electrons).

Chemical Formulas and Nomenclature

Writing Chemical Formulas

Formulas represent compounds using symbols, subscripts, and parentheses.

• Subscripts: Indicate the number of atoms of an element. If absent, the number is assumed to be $1$.

• Parentheses/Brackets: Used to enclose polyatomic ions (radicals) when more than one is present in the formula.

Steps for Formula Writing

1. Determine Symbols: e.g., Potassium ($K$), Silver ($Ag$), Aluminum ($Al$).

2. Determine Valence (Charge): e.g., $Na = +1$, $Mg = +2$, $Al = +3$.

3. Ordering: Write the positive element first, then the negative element.

4. Criss-Cross Rule: The numerical value of the valence of one ion becomes the subscript of the other. Drop the algebraic signs.

   • Example: $Al^{+3}$ and $O^{-2}$ becomes $Al_2O_3$.

Specific Formula Rules

1. Subscript of 1: Not written (e.g., $Na^{+1}Cl^{-1} \rightarrow NaCl$).

2. Equal Valences: If valences are numerically equal (e.g., both are $2$), no subscripts are written (e.g., $Ca^{+2}O^{-2} \rightarrow CaO$; $Fe^{+2}O^{-2} \rightarrow FeO$).

3. Radical Groups: If a polyatomic ion appears more than once, it must be in parentheses (e.g., $(NH_4)_2SO_4$, $Fe_2(SO_4)_3$, $Ca(OH)_2$).

Naming Binary Compounds

Compounds composed of two different kinds of atoms.

Two Non-Metals

• Name the positive ion, then the negative ion ending in "-ide."

• Greek Prefixes (Quantity):

  • $1$: mono

  • $2$: di

  • $3$: tri

  • $4$: tetra

  • $5$: penta

  • $6$: hexa

  • $7$: hepta

  • $8$: octa

  • $9$: nona

  • $10$: deca

• Examples: $CO$ (carbon monoxide), $CO_2$ (carbon dioxide), $PCl_5$ (phosphorous pentachloride), $N_2O_5$ (dinitrogen pentoxide).

Metal and Non-Metal

• Fixed Oxidation Number: Name the positive ion followed by the negative ion ending in "-ide."

  • $Al_2O_3$: aluminum oxide

  • $ZnCl_2$: zinc chloride

  • $MgH_2$: magnesium hydride

  • $Na_2S$: sodium sulfide

• Variable Oxidation Number:

  • Stock (Roman Numeral) Method: Specify the charge of the metal using Roman numerals. Example: $PbO$ is Lead (II) oxide; $PbO_2$ is Lead (IV) oxide.

  • Old Method: Use the suffix "-ous" for the lower oxidation state and "-ic" for the higher oxidation state. Example: $PbO$ is Plumbous oxide; $PbO_2$ is Plumbic oxide.

Ternary Compounds and Polyatomic Ions

Compounds composed of three or more kinds of atoms, often involving polyatomic ions.

• Common Polyatomic Ions:

  • $NH_4^+$: ammonium

  • $OH^-$: hydroxide

  • $NO_3^-$: nitrate

  • $ClO_3^-$: chlorate

  • $CO_3^{2-}$: carbonate

  • $HCO_3^-$: hydrogen carbonate (bicarbonate)

  • $SO_4^{2-}$: sulfate

  • $PO_4^{3-}$: phosphate

• Naming Examples:

  • $MnSO_4$: Manganese (II) sulfate

  • $Ca_3(PO_4)_2$: Calcium phosphate

  • $NH_4NO_3$: ammonium nitrate

  • $LiHCO_3$: Lithium bicarbonate

Acids and Bases

Acids

Hydrogen is written first. Naming depends on the anion.

• Binary Acids (H + monatomic ion): Use prefix "hydro-", the root of the anion, the suffix "-ic", and the word "acid."

  • $HCl$: hydrochloric acid

  • $H_2S$: hydrosulfuric acid

• Ternary Acids (H + polyatomic ion):

  • If the anion ends in "-ate," change the suffix to "-ic acid."

    • $H_2SO_4$ (from sulfate): sulfuric acid

  • If the anion ends in "-ite," change the suffix to "-ous acid."

    • $HBrO$ (from hypobromite): hypobromous acid

    • $H_2SO_3$ (from sulfite): sulfurous acid

Bases

A compound that produces $OH^-$ ions in aqueous solution, composed of a metal and hydroxide.

• Naming: Name the metal followed by "hydroxide."

  • $NaOH$: sodium hydroxide

  • $Ca(OH)_2$: calcium hydroxide

• Ammonia ($NH_3$): Acts as a base. When put in water, it forms Ammonium ($NH_4^+$) and releases $OH^-$ ions.

Practice Problems Summary

Formula to Name

1. $Na_3PO_4$: Sodium phosphate

2. $K_3PO_4$: Potassium phosphate

3. $Li_2CO_3$: Lithium carbonate

4. $MgCO_3$: Magnesium carbonate

5. $RbNO_3$: Rubidium nitrate

6. $Be(NO_3)_2$: Beryllium nitrate

7. $AgNO_3$: Silver nitrate

8. $Al_2(SO_4)_3$: Aluminum sulfate

9. $(NH_4)_2SO_4$: Ammonium sulfate

10. $MgF_2$: Magnesium fluoride

Name to Formula

1. Potassium fluoride: $KF$

2. Ammonium sulfate: $(NH_4)_2SO_4$

3. Magnesium iodide: $MgI_2$

4. Copper (II) sulfite: $CuSO_3$

5. Aluminum phosphate: $AlPO_4$

6. Lead (II) nitrite: $Pb(NO_2)_2$

7. Cobalt (II) selenide: $CoSe$

8. Silver cyanide: $AgCN$

9. Copper (II) bicarbonate: $Cu(HCO_3)_2$

10. Iron (II) oxide: $FeO$