Chemical Names, Formulas, and Oxidation States Notes on Composition

Significance of Chemical Formulas and Names

• Total Compounds: The number of natural and synthetic chemical compounds is currently in the millions.

• Common Naming vs. Systematic Naming:

  • Common names often provide no information regarding chemical composition and are used colloquially.

  • Examples: Calcium carbonate is known as limestone; sodium chloride is table salt; dihydrogen monoxide is water.

  • Systematic Naming: Chemists use standardized methods to describe the atomic makeup of compounds.

• Molecular Compounds: The chemical formula reveals the exact number of atoms of each element in a single molecule of the compound.

  • Example: Octane ($C_{8}H_{18}$) contains 8 carbon atoms and 18 hydrogen atoms.

  • Hydrocarbons: Molecular compounds composed solely of carbon and hydrogen.

• Ionic Compounds: Consist of a lattice of positive and negative ions held together by mutual attraction.

  • Formula Unit: The simplest ratio of the compound’s positive ions (cations) and negative ions (anions).

• Formula Interpretation Case Study: Aluminum Sulfate $Al_{2}(SO_{4})_{3}$

  • Subscript 2: Signifies 2 aluminum atoms in the formula unit.

  • Subscript 4: Signifies 4 oxygen atoms specifically within the sulfate polyatomic ion.

  • Parentheses: Surround the polyatomic anion ($SO_{4}$) to identify it as a functional unit.

  • Subscript 3 (Outside Parentheses): Refers to everything inside, giving 3 sulfate ions total, resulting in 3 sulfur atoms and 12 oxygen ($3 \times 4$) atoms.

  • Understood Subscripts: If no subscript is written, the value is understood to be 1.

Monatomic Ions and Ion Formation

• Definition: Monatomic ions are ions formed from a single atom by gaining or losing electrons to achieve noble-gas configurations (completed outermost octets).

• Cations (Positive Ions):

  • Group 1 Metals: Lose one electron to form $1+$ cations (e.g., $Na^{+}$).

  • Group 2 Metals: Lose two electrons to form $2+$ cations (e.g., $Mg^{2+}$).

  • Group 13 Metals: Typically form $3+$ cations (e.g., $Al^{3+}$).

  • Transition elements (d-block): Often form ions with variable charges, such as $2+$, $3+$, or sometimes $1+$ and $4+$.

    • Copper: $Cu^{+}$ and $Cu^{2+}$.

    • Iron and Chromium: $2+$ and $3+$ cations.

    • Vanadium: $2+$, $3+$, and $4+$ cations.

• Anions (Negative Ions):

  • Most nonmetals of Group 15, 16, and 17 gain electrons.

  • Group 15: Gain three electrons to form $3-$ anions (e.g., Nitrogen forms nitride, $N^{3-}$).

  • Group 16: Gain two electrons to form $2-$ anions (e.g., Oxygen forms oxide, $O^{2-}$).

  • Group 17 (Halogens): Gain one electron to form $1-$ anions (e.g., Fluoride, $F^{-}$).

• Exceptions and Variations:

  • Carbon and Silicon: Do not readily form ions; they typically form covalent bonds by sharing electrons.

  • Groups 14 Metals (Tin and Lead): Difficult to lose four electrons. They often lose only two electrons from their outer p orbitals ($2+$ cations, like $Sn^{2+}$ or $Pb^{2+}$), though they can also form molecular compounds using all four valence electrons.

• Naming Monatomic Ions:

  • Cations: Identified simply by the element's name (e.g., Potassium cation).

  • Anions: Drop the ending of the element name and add "-ide" (e.g., Chlorine becomes chloride).

Binary Ionic Compounds and the Stock System

• Binary Compounds: Compounds composed of atoms of only two different elements.

• Balancing Charges: In a binary ionic compound, the total number of positive charges must equal the total negative charges.

  • Example: Magnesium Bromide. Magnesium ($Mg^{2+}$) and Bromine ($Br^{-}$) combine. To balance, two bromine ions are needed: $MgBr_{2}$.

• The Crossing Over Method: A structural technique to determine subscripts.

  1. Write the symbols side by side (cation first). Example: $Al^{3+}$ and $O^{2-}$.

  2. Cross over the absolute value of the charges to become subscripts for the opposite ion: $Al_{2}O_{3}$.

  3. Check and divide subscripts by their greatest common factor to ensure the smallest whole-number ratio.

• Stock System of Nomenclature:

  • Uses Roman numerals in parentheses to indicate an ion’s charge immediately after the metal name.

  • Essential for elements forming multiple cations (e.g., d-block metals).

  • Example: $FeCl_{2}$ is Iron(II) chloride; $FeCl_{3}$ is Iron(III) chloride.

  • Silver ($Ag^{+}$), Zinc ($Zn^{2+}$), and Barium ($Ba^{2+}$) usually do not need Roman numerals as they typically form only one ion.

• Oxides of Lead Example:

  • $PbO_{2}$: Lead(IV) oxide.

  • $PbO$: Lead(II) oxide.

Polyatomic Ions and Oxyanions

• Oxyanions: Polyatomic ions that contain oxygen.

• Suffixes based on Oxygen Count:

  • -ate: Used for the ion with the greater number of oxygen atoms (e.g., Nitrate, $NO_{3}^{-}$).

  • -ite: Used for the ion with the lesser number of oxygen atoms (e.g., Nitrite, $NO_{2}^{-}$).

• Extended Series (Chlorine Example):

  • Hypochlorite ($ClO^{-}$): One less oxygen than "-ite".

  • Chlorite ($ClO_{2}^{-}$): Base ending for lesser oxygen.

  • Chlorate ($ClO_{3}^{-}$): Base ending for greater oxygen.

  • Perchlorate ($ClO_{4}^{-}$): One more oxygen than "-ate".

• Common Polyatomic Ions Checklist:

  • Ammonium: $NH_{4}^{+}$

  • Carbonate: $CO_{3}^{2-}$

  • Hydroxide: $OH^{-}$

  • Sulfate: $SO_{4}^{2-}$

  • Phosphate: $PO_{4}^{3-}$

  • Acetate: $CH_{3}COO^{-}$

  • Permanganate: $MnO_{4}^{-}$

• Special Note on Mercury(I): Exists as two $Hg^{+}$ ions joined by a covalent bond, written as $Hg_{2}^{2+}$.

Binary Molecular Compounds and Prefix System

• Prefix System: Used primarily for covalent molecular units.

• Numerical Prefixes:

  • 1: mono-

  • 2: di-

  • 3: tri-

  • 4: tetra-

  • 5: penta-

  • 6: hexa-

  • 7: hepta-

  • 8: octa-

  • 9: nona-

  • 10: deca-

• Naming Rules:

  1. Smaller group number element is listed first. If in the same group, the greater period number comes first.

  2. Give the first element a prefix only if it contributes more than one atom.

  3. The second element always has a prefix (indicating count), the root name, and the ending "-ide".

  4. Drop "o" or "a" at the end of a prefix if the element starts with a vowel (e.g., monoxide, not mono-oxide).

• Nitrogen Oxides Examples:

  • $N_{2}O$: dinitrogen monoxide.

  • $NO_{2}$: nitrogen dioxide.

  • $N_{2}O_{5}$: dinitrogen pentoxide.

• Covalent-Network Compounds: Large 3D networks of atoms (no single molecules).

  • Examples: Silicon carbide ($SiC$), Silicon dioxide ($SiO_{2}$), Trisilicon tetranitride ($Si_{3}N_{4}$).

Acids and Salts

• Definition: An acid is a molecular compound usually recognized as such when in water solution.

• Binary Acids: Consist of hydrogen and a halogen (e.g., $HCl$ is hydrochloric acid).

• Oxyacids: Contain hydrogen, oxygen, and a third element (usually a nonmetal).

  • Sulfuric acid: $H_{2}SO_{4}$.

  • Nitric acid: $HNO_{3}$.

  • Phosphoric acid: $H_{3}PO_{4}$.

• Salts: Ionic compounds composed of a cation and the anion from an acid.

  • $NaCl$ contains the anion from hydrochloric acid.

  • Hydrogen-containing anions are named with "hydrogen" or "bi-" (e.g., $HCO_{3}^{-}$ is hydrogen carbonate or bicarbonate).

Oxidation Numbers (Oxidation States)

• Definition: Assigned to atoms to indicate the general distribution of electrons in molecular compounds or polyatomic ions. They are arbitrary and do not have exact physical meaning like ionic charges.

• Rules for Assignment:

  1. Pure elements: $0$.

  2. Most electronegative element in a binary molecular compound gets the charge it would have as an anion.

  3. Fluorine: Always $-1$.

  4. Oxygen: Almost always $-2$. Exceptions: $-1$ in peroxides ($H_{2}O_{2}$) and $+2$ in $OF_{2}$.

  5. Hydrogen: $+1$ with more electronegative elements; $-1$ with metals.

  6. Neutral Compounds: Sum of all oxidation numbers must be $0$.

  7. Polyatomic Ions: Sum of all oxidation numbers must equal the ion's charge.

  8. Monatomic Ions: Oxidation number equals the ion's charge.

• Application: Used as an alternative to prefixes in the Stock system for molecular compounds (e.g., $SO_{2}$ is sulfur(IV) oxide).

Formula Mass, Molar Mass, and Percentage Composition

• Formula Mass: Sum of the average atomic masses of all atoms in a formula (unit: $u$).

  • Water ($H_{2}O$) calculation: $(2 \times 1.01) + (1 \times 16.00) = 18.02\,u$.

• Molar Mass: The mass in grams of one mole of a substance ($g/mol$). Numerically equal to formula mass.

  • Barium nitrate ($Ba(NO_{3})_{2}$) calculation: $(1 \times 137.33) + (2 \times 14.01) + (6 \times 16.00) = 261.35\,g/mol$.

• Conversion Factors:

  • $\text{moles} \times \text{molar mass (g/mol)} = \text{mass (g)}$.

  • $\text{mass (g)} \times \frac{1}{\text{molar mass (g/mol)}} = \text{amount in moles}$.

  • $\text{moles} \times 6.022 \times 10^{23} = \text{molecules/atoms}$.

• Percentage Composition: The mass percentage of each element in a compound.

  • Formula: $\text{\% element} = \frac{\text{mass of element in 1 mol of compound}}{\text{molar mass of compound}} \times 100$.

  • Example: $BaF_{2}$ is $78.32\%\,Ba$ and $21.66\%\,F$.

  • Hydrates Example: Sodium carbonate decahydrate ($Na_{2}CO_{3} \cdot 10H_{2}O$). Mass of 10 moles of water divided by total molar mass gives $62.97\%\,H_{2}O$.

Determining Chemical Formulas

• Empirical Formula: Represents the smallest whole-number mole ratio of atoms in a compound.

  • For ionic compounds, this is usually the formula unit.

  • For molecular compounds, it may differ from the actual molecular formula.

• Calculation Process:

  1. Convert percentage composition to mass composition (assume a $100\,g$ sample).

  2. Convert mass to moles by dividing by atomic mass.

  3. Divide each mole amount by the smallest mole value in the set.

  4. Round to nearest whole numbers or multiply to clear fractions (e.g., a $2.5$ ratio becomes $5$).

• Molecular Formula: The actual formula, representing the true number of atoms.

  • Relationship: $x(\text{empirical formula}) = \text{molecular formula}$.

  • Finding $x$: $x = \frac{\text{molecular formula mass}}{\text{empirical formula mass}}$.

  • Example: BH$_{3}$ (empirical) has mass $13.84\,u$. If molecular mass is $27.67\,u$, $x = 2$, making the molecular formula $B_{2}H_{6}$.

Insights: Science and Professional Applications

• Pharmacists: Health professionals who develop customized drug treatment plans, advise on side effects/interactions, and work in research for agencies and pharmaceutical companies.

• Mass Spectrometry:

  • Function: Measures mass-to-charge ratios of ions to identify unknown compounds and structures.

  • Sensitivity: Uses very small samples ($10^{-12}\,g$).

  • Key Figures: John Fenn and Koichi Tanaka won the 2002 Nobel Prize in Chemistry for developing electrospray ionization and MALDI techniques, allowing the study of large biological macromolecules like proteins and steroids.

Questions & Discussion

• Check for Understanding (Atomic Characteristics): What determines if atoms form ionic or covalent compounds? This depends on electronegativity differences and whether valence electrons are exchanged or shared.

• Discussion (Need for Naming Systems): Why are there multiple nomenclature methods? Different systems (prefixes vs. Stock system/oxidation numbers) are necessary to accommodate the vast complexity and different bonding types (ionic, molecular) within chemical compounds.

• Math Check (Percentage Composition): To find the % Nitrogen in ammonium sulfate, $(NH_{4})_{2}SO_{4}$, first calculate the total molar mass ($132.16\,g$) and divide the mass of 2 moles of nitrogen ($28.02\,g$) by that total, resulting in $21.20\%\,N$.