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Chemistry Lecture Notes - Key Vocabulary (Pages 1-3)

Mixtures and Properties

  • Heterogeneous Mixture: a mixture that is not the same throughout.
  • Homogeneous Mixture (Solution): a mixture that is the same throughout.
  • Physical Properties: can be measured or observed without changing the composition or identity of the substance
    • Examples: density, conductivity, melting point, color, hardness.
  • Chemical Properties: describe the way a substance may change or react to form other substances
    • Examples: flammability, corrosivity, reactivity.
  • Chemical reactions occur during chemical changes.

Subatomic Particles and Basic Concepts (Chapter 2)

  • Electron is the lightest particle (essentially weightless relative to others).
  • Electric current convention:
    • 1 C/s = 1 A
      • "One coulomb of electric charge flows through a conductor in one second" is equivalent to an electric current of one ampere.
    • Represented as 1\,\mathrm{C\,s^{-1}} = 1\,\mathrm{A}.
  • Periodic Table organization:
    • Elements are arranged in order of increasing number of protons (atomic number, Z).
    • Not all atoms have the same number of neutrons; isotopes exist.
    • Hydrogen has 1 proton and 0 neutrons; helium has 2 protons and 2 neutrons.
  • Isotopes:
    • Atoms with the same number of protons (Z) but different numbers of neutrons (N).
    • Mass number A = number of protons + number of neutrons: A = p + n\; (p = \text{protons},\; n = \text{neutrons})
  • Atomic Symbols:
    • Mass number: superscript A; Atomic number: subscript Z.
    • Example: for a neutral titanium example: $^{48}_{22}\mathrm{Ti}$
    • Protons = 22
    • Neutrons = 26 (since A - Z = 48 - 22 = 26)
    • Electrons in neutral atom = 22 (same as protons)
  • Ions:
    • When neutral atoms gain or lose electrons, they become ions.
    • Cations (positive) typically formed by metals (e.g., Mg^{2+} loses 2 electrons).
    • Anions (negative) typically formed by non-metals (e.g., O^{2-} gains 2 electrons).
    • Example: $^{51}_{23}\mathrm{V^{4+}}$ has Protons = 23, Neutrons = 28 (since A = 51\;\Rightarrow\; N = A - Z = 51 - 23 = 28). Electrons = Z - \text{charge} = 23 - 4 = 19.
  • Example for earlier elements:
    • Nickel (II) Cyanide: Ni^{2+} CN^{-} = Ni(CN)_{2}.

Isotopes, Mass, and Atomic Weights

  • Isotopes: atoms with the same Z but different N; some isotopes are more abundant than others.
  • Mass Spectrometry:
    • Ionized isotopes are separated according to their mass-to-charge ratio \frac{m}{z}.
    • Relative abundances can be calculated from the mass spectrum.
  • Atomic Weights (amu):
    • Atomic weight is the weighted average of all isotopes of an element.
    • \mathrm{amu} = \sumi (\text{abundance}i \times \text{isotopic\ mass}_i)
    • The result depends on isotope abundances.
    • Important: use abundances as fractions (not percentages) when weighting; do not inappropriately add masses with differing abundances.

Periodic Table Fundamentals

  • Periods: rows.
  • Groups: columns.
  • Notable groups:
    • 1A: Alkali metals (excluding H)
    • 2A: Alkaline earth metals
    • 6A: Chalcogens
    • 7A: Halogens
    • 8A: Noble gases
  • Other classifications: metals, metalloids, non-metals.

Chemical Compounds and Formulas

  • Chemical compounds form from ions or covalent bonds:
    • Ionic compounds: formed between cations and anions (e.g., metal cation + nonmetal anion).
    • Covalent compounds: formed between nonmetal atoms (nonmetal + nonmetal).
  • Chemical formulas:
    • Empirical formula: simplest whole-number ratio of atoms in a compound; ionic compounds always have empirical formulas.
    • Molecular formula: exact number of atoms of each element in a molecule; covalent compounds have molecular formulas.
    • Structural formula: exact connectivity of atoms.

Naming Oxyanions and Ionic Compounds

  • Oxyanions naming:
    • -ate indicates more oxygen than its -ite counterpart.
    • Example: Sulfate (SO4^{2-})
    • -ite indicates fewer oxygens than the corresponding -ate.
    • Example: Sulfite (SO3^{2-})
  • When the anion is an element, change its ending to -ide; when the anion is a polyatomic ion, use its name as is:
    • NaNO2 → Sodium nitrite
    • NaF → Sodium fluoride
    • FeCl2 → Iron(II) chloride
    • Fe(NO2)3 → Iron(III) nitrite
    • (NH4)2S from NH4^{+} and S^{2-} ions
  • Balancing ions to form compounds:
    • Combine charges to achieve a neutral compound
    • Swap and multiply as needed to obtain the smallest whole-number ratio.
    • Rule of thumb: determine the smallest integers that balance the total positive and negative charges.

Covalent Compounds and Acids

  • Nickel(II) Cyanide example:
    • Ni^{2+} and CN^{-} combine to form Ni(CN)_{2}.
  • Naming covalent compounds:
    • Acids have H^{+} and an anion.
    • Examples:
    • H{2}CO{3} → carbonic acid
    • Acids are hydrogen-containing covalently bonded compounds that give off H^{+} in water.
  • Naming acids (binary and oxyacids):
    • Binary acids (hydrogen + nonmetal): replace the suffix of the anion with -ic and add the word acid, using hydro- prefix for the nonmetal element:
    • HF → hydrofluoric acid
    • HCl → hydrochloric acid (note: chlorine-based example; the hydro-prefix is used for binary acids of halogens; historically HCl is named hydrochloric acid, not hydro chloride; the general pattern is hydro + nonmetal root + ic acid for many halogen binaries)
    • Oxyacids (contain oxygen):
    • -ate → change to -ic acid
      • NO{3}^{-} (nitrate) → HNO{3} (nitric acid)
    • -ite → change to -ous acid
      • NO{2}^{-} (nitrite) → HNO{2} (nitrous acid)
  • Naming binary covalent bonds:
    • Use prefixes mono-, di-, tri-, etc. to denote the number of atoms of each element in the compound.
    • The first element listed typically does not receive the prefix mono- unless necessary for clarity.
    • Examples:
    • CO_{2} → Carbon dioxide
    • N{2}O{5} → Dinitrogen pentoxide
  • Organic vs Inorganic compounds:
    • Inorganic: typically nonmetal + nonmetal or ionic compounds; not containing carbon-hydrogen frameworks as a rule.
    • Organic: primarily carbon and hydrogen, often with other elements; formula general form Cn Hm with other heteroatoms.
    • Example: C{3}H{8} → propane

Additional Notes and Clarifications

  • The transcript contains some typographical errors (e.g., “disclude” and mixed phrasing). The notes above reflect corrected and standard chemical conventions where applicable while preserving the original examples.
  • Key equations and notations to remember:
    • 1\,\mathrm{C\,s^{-1}} = 1\,\mathrm{A}
    • A = p + n, \quad Z = p
    • ^{A}_{Z}\mathrm{X} notation for isotopes (A = mass number, Z = atomic number, X = element symbol)
    • \frac{m}{z} for mass spectrometry data
    • \mathrm{amu} = \sumi (\text{abundance}i \times \text{isotopic mass}_i)

Quick Reference Examples from Transcript

  • Example isotope notation: ^{51}_{23}\mathrm{V^{4+}}
    • Protons (Z) = 23
    • Neutrons = 28 (since A = 51, N = A - Z = 28)
    • Electrons = 19 (Z - charge = 23 - 4)
  • Example ionic compound and oxidation state notation:
    • \mathrm{FeCl_2} → Iron(II) chloride
    • \mathrm{Fe(NO2)3} → Iron(III) nitrite
  • Example polyatomic ion salt:
    • \mathrm{(NH4)2S} (ammonium sulfide)
  • Example acid naming patterns:
    • Binary acid: \text{HF} \rightarrow \text{hydrofluoric acid}
    • Oxyacid: \text{HNO_3} \rightarrow \text{nitric acid} (ate form)
    • Oxyacid (ite form): \text{HNO_2} \rightarrow \text{nitrous acid}
  • Example organic formula:
    • \mathrm{C3H8} → Propane