Chemistry and Chemical Reactivity – Chapter 1 and Ionic/Covalent Compounds: Study Notes

Ionic and Covalent Compounds: Naming, Ions, and Formulas

  • Ionic compounds are formed from metals and nonmetals and consist of ions (atoms or groups with positive or negative charge). Common example: table salt, NaCl (sodium chloride).
  • These compounds are often referred to as salts.
  • Types of bonding/class:
    • Nonmetal + Nonmetal → Covalent (Molecular) compounds
    • Metal + Nonmetal → Ionic compounds
Naming Monatomic Ions
  • Cation: name of the element + the word ion
    • Example: Na⁺ → sodium ion
  • Anion: name of the element with an -ide ending
    • Example: Cl⁻ → chloride ion
Metal Ions Names (1 of 2) and Transition Metals
  • All metal cations are named after their representative element followed by the word “ion.”
  • For transition metals, the metal can have more than one charge. The charge is determined by the compound's formula.
    • Example reasoning: FeCl₂ contains chloride with −1 charge; Fe must be +2 to balance: FeCl₂ → Fe²⁺ and 2 Cl⁻. So oxidation state of Fe is +2.
    • Example: FeCl₃ contains 3 Cl⁻; Fe must be +3 to balance: FeCl₃ → Fe³⁺. Oxidation state of Fe is +3.
Metal Ions Names (2 of 2) – Transition Elements
  • Transition elements often have multiple possible charges; each must be identified in the name if needed to specify charge.
  • Examples summarized:
    • Cr can be Cr²⁺, Cr³⁺, Cr⁶⁺ → chromium(II), chromium(III), chromium(VI)
    • Fe can be Fe²⁺ (iron(II), ferrous) or Fe³⁺ (iron(III), ferric)
  • Some nontransition metals also have multiple charges (e.g., Tin and Lead) described below.
Other Metals with Multiple Charges
  • Tin (Sn): +2 (stannous, Sn²⁺) and +4 (stannic, Sn⁴⁺); +4 more common
    • Tin(II) chloride: SnCl₂; Tin(IV) chloride: SnCl₄
    • Tin(II) oxide: SnO; Tin(IV) oxide: SnO₂
  • Lead (Pb): +2 (plumbous) and +4 (plumbic); +2 more common
Nonmetal Ions Names
  • Anions from nonmetals: name ends with -ide (e.g., fluoride, chloride, oxide, sulfide).
Polyatomic Ions
  • A polyatomic ion is a group of atoms that acts as a single ion with a net charge (can be anions or cations).
  • Polyatomic ions are often used in ionic compounds; their names do not get changed when forming salts (e.g., NaNO₃ = sodium nitrate).
  • Hydronium ion: H₃O⁺; note that many oxyanion series end in -ate or -ite (e.g., sulfate SO₄²⁻, sulfite SO₃²⁻).
Naming Type (Ⅱ) or B Metals (Transition Ions)
  • When metals (especially transition metals) form different oxidation states, use Roman numerals to indicate the charge on the metal.
    • Examples:
    • iron(II) chloride: FeCl₂
    • iron(III) chloride: FeCl₃
  • Classic examples of old vs modern naming:
    • Fe²⁺ → iron(II) (ferrous)
    • Fe³⁺ → iron(III) (ferric)
    • Cu⁺ → copper(I) (cuprous)
    • Cu²⁺ → copper(II) (cupric)
    • Pb²⁺ → lead(II) (plumbous)
    • Pb⁴⁺ → lead(IV) (plumbic)
    • Hg₂²⁺ → mercurous; Hg²⁺ → mercuric
Atomic Numbers and Common Charges (Selected Elements)
  • Ti (22): +2, +3, +4 (Yes, transition metal)
  • V (23): +2, +3, +4, +5 (Yes)
  • Cr (24): +2, +3, +6 (Yes)
  • Mn (25): +2, +3, +4, +6, +7 (Yes)
  • Fe (26): +2, +3 (Yes)
  • Co (27): +2, +3 (Yes)
  • Ni (28): +2, +3 (Yes)
  • Cu (29): +1, +2 (Yes)
  • Ag (47): +1 (rarely +2) (Yes)
  • Sn (50): +2, +4 (No common charge emphasis given here as transition analog; classic nomenclature used)
  • Sb (51): +3, +5 (No common charge emphasis here)
  • Au (79): +1, +3 (Yes)
  • Hg (80): +1, +2 (No)
  • Pb (82): +2, +4 (No)
  • Bi (83): +3, +5 (No)
IUPAC Naming Rules for Ionic Compounds
  • Name the cation first (metal/positive ion) using the element’s name.
    • Na⁺ → sodium
  • Name the anion second, changing the ending to -ide for monoatomic ions; polyatomic ions keep their common name (no -ide change).
    • Cl⁻ → chloride
    • NO₃⁻ → nitrate
  • If the metal can have more than one charge (especially transition metals), show the charge with Roman numerals in parentheses after the cation name (crossover rule ensures net neutral charge):
    • FeCl₂ → iron(II) chloride
    • FeCl₃ → iron(III) chloride
  • Polyatomic ions retain their names: NaNO₃ → sodium nitrate, CaSO₄ → calcium sulfate
Formula Balancing and Neutral Compounds
  • Ionic compounds are neutral overall: total positive charge equals total negative charge.
  • For monatomic ions: + charge magnitude equals − charge magnitude in the compound.
  • For polyatomic ions: use parentheses as needed to balance charges in the formula.
Formulas of Ionic Compounds: Examples
  • Sodium sulfide: Na₂S (the 1 subscript is understood/omitted for Na⁺; S²⁻ requires two Na⁺ to balance)
  • Calcium hydroxide: Ca(OH)₂ (Ca²⁺ balances two OH⁻ groups; note use of parentheses for polyatomic OH⁻)

Chapter 1: Basic Concepts of Chemistry

Chemistry and Its Methods
  • Chemistry is the study of matter and its properties, composition, and changes.
  • For any observation, we classify it as qualitative or quantitative:
    • Qualitative observations: no numbers; color, appearance, odor, texture, etc.
    • Quantitative observations: involve measurements with numbers and units.
Hypothesis, Law, Theory
  • Hypothesis: tentative explanation/prediction derived from observations.
  • Law: concise statement of a relationship that is universally true under the stated conditions.
  • Theory: well-tested, unifying explanation of a body of facts and laws; can suggest new hypotheses for testing.
Qualitative vs Quantitative Observations (Examples)
  • Qualitative: color, odor, state (solid/liquid/gas), hot/cold, etc.
  • Quantitative: mass, volume, temperature, density, etc.
Classification of Matter
  • States of matter: solid, liquid, gas.
  • Composition: element, compound, mixture.
  • Properties: physical properties (color, odor, density, melting/boiling points, solubility, conductivity, etc.) vs chemical properties (reactivity, combustibility).
  • Physical changes do not alter composition; chemical changes do (e.g., iron rusting, combustion).
States of Matter (Properties)
  • Solid: definite shape and volume; low kinetic energy.
  • Liquid: definite volume, indefinite shape; intermediate kinetic energy.
  • Gas: neither definite shape nor volume; high kinetic energy.
Energy: Kinetic and Potential
  • Kinetic energy associated with motion (thermal energy, macroscopic motion, electron movement in a conductor, wave motion).
  • Potential energy arises from position (gravitational, stretched springs, chemical energy in bonds, electrostatic energy, nuclear energy).
Macroscopic vs Submicroscopic (Particulate) Levels
  • Macroscopic: observable phenomena and measurements (experiments and data).
  • Submicroscopic/Particulate: atoms and molecules that underlie macroscopic observations.
Scientific Model: Atoms and Molecules
  • Atom: smallest unit of an element that maintains element properties.
  • Molecule: smallest unit of a compound that retains its properties; may contain more than one atom and element.
  • Models help visualize how atoms and molecules exist and interact.
Classifying Matter: Pure Substances and Mixtures
  • Pure substance: has well-defined physical and chemical properties; elements or compounds.
  • Elements: substances consisting of a single type of atom; cannot be decomposed by ordinary chemical means.
  • Compounds: composed of two or more elements in fixed ratios; can be broken down by chemical reactions.
  • Periodic Table organizes elements; 118 elements currently recognized.
Mixtures: Homogeneous vs Heterogeneous
  • Homogeneous mixture: uniform composition at the molecular level; no visible boundaries.
  • Heterogeneous mixture: nonuniform composition; components are visually distinguishable.
  • Both yield pure substances upon separation.
  • Separation methods include decantation, filtration, distillation, chromatography, etc.
Physical Properties and Changes
  • Physical properties can be observed without changing composition: color, density, melting/boiling points, solubility, electrical conductivity, etc.
  • Density is a key physical property that can vary with temperature.
  • Intensive properties do not depend on the amount of substance (e.g., density, temperature).
  • Extensive properties depend on the amount of substance (e.g., mass, volume).
Chemical Properties and Changes
  • Chemical properties describe a substance’s potential to undergo chemical change (e.g., combustibility of wood).
  • Chemical change results in new substances with different properties (e.g., iron reacting with oxygen to form rust).
Subatomic Particles and Atomic Structure
  • Atoms are composed of protons (p+), neutrons (n⁰), and electrons (e⁻).
  • Protons carry positive charge; electrons carry negative charge; neutrons are neutral.
  • Atoms are overall electrically neutral: number of protons equals number of electrons.
  • Mass number A = number of protons + number of neutrons.
  • Atomic number Z = number of protons (identifies the element).
  • Isotopes: atoms of the same element (same Z) with different A due to differing neutrons.
  • Isotopes notation: ^A_Z X, where X is the element symbol.
  • Atomic weights on the periodic table are weighted averages of isotopic masses.
Atomic Mass, Isotopes, and Abundances
  • Masses: proton ≈ 1.007 u; neutron ≈ 1.009 u; electron ≈ 5×10⁻⁴ u (negligible for mass balance).
  • Hydrogen example: isotopes include protium (¹H), deuterium (²H), tritium (³H); abundances give atomic weight.
  • Atomic weight example: Hydrogen = 1.008 amu due to isotope abundances.
The Periodic Table: Organization and Regions
  • Elements arranged by increasing atomic number (Z).
  • Periods (rows) and Groups/Columns (families) with shared properties.
  • Major families: metals, nonmetals, metalloids; halogens; noble gases; alkali metals.
  • Electron configurations underpin periodic trends (s, p, d, f blocks).
  • Lanthanide and Actinide series discussed as separate blocks.
Molecules, Formulas, and Nomenclature
  • Molecules: smallest unit of a compound retaining chemical properties; can be binary (two elements) or multi-element.
  • Formulas illustrate the types and numbers of atoms:
    • Molecular Formulas: show numbers of each type of atom in a molecule.
    • Empirical Formulas: simplest whole-number ratio of elements in a compound.
    • Structural Formulas: show how atoms are connected (bonds).
    • Condensed Formulas: groupings of atoms (e.g., CH₃OH).
    • Empirical formula for C4H10 can be written as C2H5 (simplest ratio).
Naming Molecular Compounds (Nonmetal + Nonmetal)
  • Use prefixes to indicate the number of each atom (mono-, di-, tri-, tetra-, penta-, hex-a-, etc.).
    • Example: CO₂ = carbon dioxide; N₂O₅ = dinitrogen pentoxide
  • The first element keeps its name (no prefix for the first element).
    • Example: CO = carbon monoxide (not monocarbon monoxide)
  • The second element ends with -ide.
    • Example: PCl₃ = phosphorus trichloride
  • When multiple forms exist, prefixes specify numbers to distinguish compounds with the same elements (e.g., CO vs CO₂).
Ionic Compounds: Review and Formulas
  • Ionic compounds are composed of ions; examples include NaCl, CaSO₄, Na₂S, Ca(OH)₂.
  • Balance charges to obtain neutral compounds; often require Roman numerals for transition metals in the cation.
Common Polyatomic Ions (Recap)
  • Nitrate: NO₃⁻
  • Sulfate: SO₄²⁻
  • Chlorate: ClO₃⁻
  • Hydronium: H₃O⁺
  • Acetate: C₂H₃O₂⁻ (not shown in slides but commonly encountered)

Calculations and Formulas: Key Formulas to Remember

  • Avogadro's Number

    • NA=6.022imes1023extparticles/molN_A = 6.022 imes 10^{23} ext{ particles/mol}

  • Molar Mass and Moles

    • Molar mass: the mass of one mole of a substance, in g/mol
    • ext{Moles} = rac{ ext{Mass (g)}}{ ext{Molar Mass (g/mol)}}
    • extMass(g)=extMolesimesextMolarMass(g/mol)ext{Mass (g)} = ext{Moles} imes ext{Molar Mass (g/mol)}

  • Particles from Moles

    • extParticles=extMolesimesNA=extMolesimes6.022imes1023ext{Particles} = ext{Moles} imes N_A = ext{Moles} imes 6.022 imes 10^{23}

  • Percent Composition

    • For an element X in a compound: ext{%X} = rac{nX imes MX}{M_{ ext{formula}}} imes 100eta
    • Where nX is the number of atoms of element X in the formula, MX is the atomic mass of X, and M_{ ext{formula}} is the formula mass of the compound.

  • Example: Ca(OH)₂

    • FW (formula weight) = 74.1 g/mol
    • %Ca = (40.0 g Ca per 74.1 g) × 100% = 54.0%
    • %O and %H calculated similarly (O: 16.0 × 2 = 32.0; H: 1.008 × 2 = 2.016; then divide by 74.1 and multiply by 100%)

  • Empirical Formula from Percent Composition

    • Step 1: assume 100 g of compound; convert each element’s mass to moles
    • Step 2: divide by the smallest mole value to obtain the mole ratio
    • Step 3: multiply all ratios by the same factor to obtain whole numbers
    • Step 4: decide whether to round or multiply to obtain whole numbers (rounded if within ±0.05 of a whole number; multiply by small integers like 2, 3, 4, or 5 if fractions are near common fractions)

  • Deriving Molecular Formula from Empirical Formula

    • If the empirical formula is known and molar mass is measured:
    • Compute n = rac{M{ ext{molecular}}}{M{ ext{empirical}}}
    • If n is an integer, molecular formula = (empirical formula)ⁿ
    • Example: Empirical C₄H₁₀O mass = 74.12 g/mol; if molecular molar mass = 222.1 g/mol, then n = rac{222.1}{74.12}
      ightarrow n ext{ is approximately } 3, thus molecular formula = C₁₂H₃₀O₃.

  • Example Worked Steps (Empirical from Slide Example 1)

    • Magnetite: Fe₃O₄ with Fe mass 72.36% and O mass 27.64%
    • Step 1: Convert to grams: Fe = 72.36 g, O = 27.64 g (assuming 100 g sample)
    • Step 2: Convert to moles:
    • n{ ext{Fe}} = rac{72.36}{55.85} ext{ mol} \ n{ ext{O}} = rac{27.64}{16.00} ext{ mol}
    • Step 3: Determine simplest whole-number ratio by dividing by smallest moles; adjust to whole numbers by multiplying as needed (e.g., multiply Fe:O ratio to get Fe₃O₄)

  • Additional Example: Calculation Type 1 (Conversions)

    • Grams ⇄ Moles ⇄ Particles: Follow the conversion flow: Grams ⇄ Moles ⇄ Particles.
    • For grams to moles: use molar mass; for moles to particles: multiply by NA=6.022imes1023N_A = 6.022 imes 10^{23}.

  • Example Problem (Moles to Mg atoms)

    • Given Mg mass, compute moles using Mg’s molar mass, then convert to number of Mg atoms using NAN_A.

Quick Reference: Key Concepts to Memorize

  • Ionic vs Covalent bonding: metals typically form cations; nonmetals form anions.
  • Roman numerals in metal names indicate oxidation state (transition/metals with multiple charges).
  • Polyatomic ions retain their names in ionic formulas (no -ide changes).
  • Balance the total positive and negative charges to obtain neutral ionic formulas.
  • Isotopes share Z (same protons) but have different A (mass numbers) due to different neutron counts.
  • Atomic weight on the periodic table is a weighted average of isotope masses, reflecting natural abundances.
  • The empirical formula represents the simplest whole-number ratio of elements; the molecular formula is a whole-number multiple of the empirical formula.
  • The mole is the bridge between atomic scale and lab scale; Avogadro's number connects moles to particles.

Notes:

  • When presenting chemical formulas in your notes, you can express them inline as extNa2extSext{Na}_2 ext{S} or as plain text; use the LaTeX formatting for calculations and formula masses as shown above where appropriate.
  • For transition metal compounds, always verify charge balance and use Roman numerals in the cation name when the metal has multiple oxidation states.