Notes on Chapter 3: Molecules and Compounds (Key Concepts and Details)

Rocket Fuel

• Hydrogen, Oxygen, Water: Selected properties
  • Hydrogen
  • Boiling Point: $-253\,^{\circ}\mathrm{C}$
  • State at room temperature: Gas
  • Flammability: Explosive
  • Oxygen
  • Boiling Point: $-183\,^{\circ}\mathrm{C}$
  • State at room temperature: Gas
  • Flammability: Necessary for combustion (supports combustion)
  • Water
  • Boiling Point: $100\,^{\circ}\mathrm{C}$
  • State at room temperature: Liquid
  • Flammability: Used to extinguish flame
• Rocket Fuel context: mixtures and compounds
  • Hydrogen and Oxygen mixture
  • Can have any ratio of hydrogen to oxygen
  • Water (a compound)
  • Water molecules have a fixed ratio of atoms: $2\,\mathrm{H} : 1\,\mathrm{O}$
• Reference prompt video: Reactions with Oxygen (see Course Videos)

Chemical Bonds

• Two general types of bonds that hold atoms together:
  • Ionic Bond
  • Attraction between a cation (positive charge) and an anion (negative charge)
  • Covalent Bond
  • Sharing of a pair of electrons, with one electron coming from each atom
• These are different, but part of a range of bonding – based on sharing electrons

Ionic Bonding in Solids

• In a solid, ions of opposite charge are arranged in a 3D lattice (lattice structure)

Covalent Bonds

• Covalent bonds share electrons
  • Atoms each share one electron to create a bonding pair
  • The bonding pair is energetically stable

Molecular Formulae

• Molecular formulae: Written description of the elements making up a compound, and some structural information
• Empirical formula
  • Simplest whole-number ratio of atoms of each element in the compound
  • Example: for hydrogen peroxide, empirical formula is $HO$
• Molecular formulae
  • Actual number of atoms of each element in the compound
  • Example: for hydrogen peroxide, molecular formula is $H<em>2O</em>2$
• Structural formulae
  • Shows the actual number of atoms and how they are bonded together
• Increasing information from empirical to molecular to structural representations

Examples of Molecular Formulae

• C6H12: hexene
  • Empirical formula: $CH_2$
• C5H10: pentene
  • Empirical formula: $CH_2$
• C4H8: butene
  • Empirical formula: $CH_2$
• Methane
  • Empirical formula: $CH<em>4$; Molecular formula: $CH</em>4$ (same for methane)
• The more atoms in a molecular formula, the larger the number of possible compounds
• Example: $C<em>6H</em>{12}O_6$ can have at least $1,587$ compounds
  • Some sample representations:
  • $CH<em>3CH</em>2CH<em>2CH=CHCH</em>3$
  • $CH<em>3CH</em>2CH=CHCH_3$
  • $CH<em>3CH=CHCH</em>3$
• Source note: PubChem example for $C<em>6H</em>{12}O_6$

Structural Formula with Wedge Bonds

• D-Glucose (2D structure) vs chemical structure depiction
• Shows how bonds and atoms are arranged in 2D form

Molecular Models of Structure

• To visualize molecules in 3D with realistic atom sizes, models are used
  • 3D molecular models allow better understanding of spatial arrangement

Different Views for Different Needs (Table Overview)

• Name of compound
• Empirical formula
• Molecular formula
• Structural formula
• Ball-and-stick model
• Space-filling model
• Examples:
  • Benzene: empirical formula $CH$; molecular formula $C<em>6H</em>6$; structural formula shows a six-member ring with alternating single and double bonds; ball-and-stick model; space-filling model
  • Acetylene: empirical formula $CH$; molecular formula $C<em>2H</em>2$; space-filling model described
  • Glucose: empirical formula $CH<em>2O$; molecular formula $C</em>6H<em>{12}O</em>6$; structural formula shows a vertical chain with various attachments; ball-and-stick and space-filling models described
  • Ammonia: formula $NH_3$; ball-and-stick and space-filling models described

Compounds

• Understanding compounds by type:
  • Molecular compounds
  • Typically composed of two or more covalently bonded non-metals
  • Basic units are molecules
  • Examples:
    • Water: $H_2O$ molecules
    • Dry ice: $CO_2$ molecules
    • Propane: $C<em>3H</em>8$ molecules
  • Ionic compounds
  • Composed of cations (usually metals) and anions (usually nonmetals) bound by ionic bonds
  • Basic unit: formula unit (smallest electrically neutral group of ions)
  • Example: table salt: $NaCl$, composed of $Na^+$ and $Cl^-$ in a 1:1 ratio

Naming Ionic Compounds

• Cations (positive ions) and anions (negative ions)
• Cations typically metals; anions from nonmetals
• General rules:
  • Metals form cations
  • Halides, oxygen, sulfur form anions
  • Polyatomic cations and anions contain more than one type of atom
  • Oxyanions contain one or more oxygen atoms
• Writing formulas:
  • Cation is listed before anion
  • The formula for any polyatomic ion is written as a unit
  • Polyatomic ions are placed in parentheses with a subscript if there are two or more in the formula (e.g., $NaOH$, $Ba(OH)_2$)
• Recognizing ionic compounds:
  • Contains a metal from Group I or Group II, or one of the polyatomic ions

Visualizing Ionic Compounds

• Example diagram (conceptual):
  • Sodium atom (Na) loses an electron to become Na^+; Chlorine atom (Cl) gains an electron to become Cl^-; shows electrostatic attraction between oppositely charged ions
• Notation in diagrams often uses charge counts (e.g., 11p+, 11e) to illustrate electron transfer

Naming Cations

• Cations (positive ions) typically metals
• Names:
  • Monatomic ions (one atom) share the element name: Lithium ion (Li^+), Calcium ion (Ca^{2+})
  • Cations can have multiple charges; indicate with Roman numerals in parentheses: Iron(II) (Fe^{2+}), Iron(III) (Fe^{3+})
• Some older name forms: -ous and -ic designations for different charges
  • Iron(II) -> Ferrous ion
  • Iron(III) -> Ferric ion
• Polyatomic ions with positive charges also exist (e.g., $NH<em>4^+$ ammonium, $H</em>3O^+$ hydronium)

Common Cations

• Table 2.4 (highlights): Common Cations
• 1+ charge: $H^+$ (hydrogen ion), $Li^+$ (lithium ion), $NH_4^+$ (ammonium ion), $Na^+$ (sodium ion), $Cu^+$ (copper(I) or cuprous), $K^+$ (potassium ion), $Cs^+$ (cesium ion)
• 2+ charge: $Ag^+$ (silver ion), $Mg^{2+}$ (magnesium ion), $Ca^{2+}$ (calcium ion), $Sr^{2+}$ (strontium ion), $Ba^{2+}$ (barium ion), $Co^{2+}$ (cobalt(II) or cobaltous), $Cu^{2+}$ (copper(II) or cupric), $Fe^{2+}$ (iron(II) or ferrous), $Mn^{2+}$ (manganese(II) or manganous), $Zn^{2+}$ (zinc), $Cd^{2+}$ (cadmium), $Hg_2^{2+}$ (mercury(I) or mercurous)
• 3+ charge: $Al^{3+}$ (aluminum), $Cr^{3+}$ (chromium(III) or chromic), $Fe^{3+}$ (iron(III) or ferric)
• Note: Bolded ions are the ones used most often in this course

Naming Anions

• Monatomic anions end in -ide (e.g., Cl^-, S^{2-}, O^{2-})
• Polyatomic anions with oxygen (oxyanions) end in -ate or -ite
  • Examples: $NO<em>3^-$ nitr ate, $NO</em>2^-$ nitrite, $SO<em>4^{2-}$ sulfate, $SO</em>3^{2-}$ sulfite

Common Anions

• Table 2.5 (highlights): Common Anions
• 1−: $H^-$ (hydride), $F^-$ (fluoride), $ClO<em>3^-$ (chlorate), $Cl^-$ (chloride), $Br^-$ (bromide), $NO</em>3^-$ (nitrate), $I^-$ (iodide), $CN^-$ (cyanide), $OH^-$ (hydroxide), $O_2^-$ (oxide)
• 2−: $O^{2-}$ (oxide), $CO<em>3^{2-}$ (carbonate), $O</em>2^{2-}$ (peroxide), $CrO<em>4^{2-}$ (chromate), $Cr</em>2O<em>7^{2-}$ (dichromate), $SO</em>4^{2-}$ (sulfate), $PO_4^{3-}$ (phosphate)
• 3−: $N^{3-}$ (nitride), etc.
• Note: Bolded ions are the most commonly used ones in this course

Charges on Elemental Ions (Overview)

• Illustration of how groups 1A–3A, transition metals, metalloids, and nonmetals relate to common ion charges
• General idea: alkali and alkaline earth metals form simple cations; transition metals can have multiple charges; nonmetals form anions or oxyanions depending on composition

Naming Acids

• Acids are named based on the accompanying anion
• Monoatomic anions:
  • Ide: add H+, prefix hydro-, end with -ic acid (e.g., chloride → hydrochloric acid; HCl → hydrochloric acid)
• Oxyanions:
  • -ate: end with -ic acid (e.g., chlorate NO3− → chloric acid HClO3)
  • -ite: end with -ous acid (e.g., chlorite NO2− → chlorous acid HClO2)
• Hypo- and per- cases (oxygen-containing anions):
  • Hypochlorite (ClO−) → hypochlorous acid (HClO)
  • Chlorite (ClO2−) → chlorous acid (HClO2)
  • Chlorate (ClO3−) → chloric acid (HClO3)
  • Perchlorate (ClO4−) → perchloric acid (HClO4)
• Note: Acids must be in aqueous solution

Naming Molecular Compounds

• Also called Binary Molecular Compounds
• Formed between two non-metal elements
• The ratio of atoms of each element can vary, so a systematic naming method is needed
• Guidelines for naming two-element compounds:
  • The element that is farthest left in the periodic table goes first
  • If in the same group, the lower element goes first
  • The name of the second element ends with -ide
  • Prefixes indicate the number of atoms of each element (do not use mono for the first element)

Naming Prefixes

• Prefix rules:
  • Mono-, Di-, Tri-, Tetra-, Penta-, Hexa-, Hepta-, Octa-, Nona-, Deca-
  • When a prefix ends in -a or -o and the name begins with a vowel, the a or o is dropped (e.g., carbon dioxide not monooxide carbon dioxide, etc.)

Inorganic Naming Flowchart (Overview)

• IONIC (Metal and nonmetal)
  • Metal forms one type of ion only: name of cation + base name of anion (nonmetal) + -ide
  • Example: CaI2 → calcium iodide
• IONIC (Metal with multiple cation charges)
  • Name of cation + charge of cation in Roman numerals in parentheses + name of anion
  • Example: FeCl3 → iron(III) chloride
• MOLECULAR (Nonmetals only)
• ACIDS*
  • H and one or more nonmetals
  • Binary acids: hydro + base name of nonmetal + -ic acid (e.g., HCl → hydrochloric acid)
  • Oxyacids: base name of oxyanion + -ic or -ous acid depending on the oxyanion
  • Example: P2O5 → diphosphorus pentoxide; H2SO3 → sulfurous acid
  • *Acids must be in aqueous solution

Calculating Empirical Formulae

• Analysis steps:
  • Measure the mass of each element in the compound
  • Express as % of the total mass; assume a 100 g sample for convenience
  • Calculate the number of moles of each element
  • Divide by the smallest number of moles to obtain mole ratios
  • Multiply through by a small integer to get whole numbers (ratios may be inexact)

Calculating Formula Mass

• Definition: The mass of an individual molecule or formula unit; also called molecular mass or molecular weight
• It is the sum of the masses of the atoms in a single molecule or formula unit
• Example principle: whole = sum of the parts

Calculating Molar Mass

• Molar mass: mass, in grams, of 1 mole of its molecules or formula units; numerically equal to the formula mass with units of g/mol
• Conceptual example: 1 mole of $\mathrm{H_2O}$ contains 2 moles of H and 1 mole of O

Molar Mass and Counting Molecules

• Use molar mass in combination with Avogadro’s number to determine the number of atoms in a mass of a substance
• Procedure:
  • Convert mass to moles using molar mass
  • Convert moles to number of molecules using Avogadro’s number
• Avogadro’s number is used to relate moles to the actual count of molecules (value not listed in the transcript but used conceptually here)

Quick Reference Formulas

• Empirical formula from mass data: ratio method as described above
• Structural, molecular, and empirical formula distinctions:
  • Empirical: simplest whole-number ratio
  • Molecular: actual number of atoms
  • Structural: shows bonding and arrangement

Quick Reference Examples in Context

• Water: empirical formula $HO$; molecular formula $H_2O$
• Hydrogen peroxide: empirical formula $HO$; molecular formula $H<em>2O</em>2$
• Methane: empirical and molecular both $CH_4$
• Glucose: empirical $CH<em>2O$; molecular $C</em>6H<em>{12}O</em>6$; structural depiction shows a chain with carbon, hydrogen, and hydroxyl groups
• Benzene: empirical $CH$; molecular $C<em>6H</em>6$; ring structure with alternating single/double bonds; ball-and-stick and space-filling models available
• Acetylene: empirical $CH$; molecular $C<em>2H</em>2$; linear structure
• Ammonia: $NH_3$; common ball-and-stick and space-filling models

Notes on Important Connections

• Molecular and ionic concepts tie to larger themes in chemistry: bonding types affect properties like boiling point, flammability, and conductivity
• Naming conventions encode structural information: oxidation states (via Roman numerals), presence of polyatomic ions, and the distinction between ionic, molecular, and acidic compounds
• Empirical vs molecular formulas illustrate how composition relates to actual molecular structure and potential isomeric diversity
• Calculation techniques (empirical formula, molar mass, and mole-to-molecule conversions) are foundational for quantitative chemistry and stoichiometry