Chapter 2 Notes: Atoms, Atomic Theory, and Atomic Structure

Chapter 2: Atoms, the Periodic Table, and Compound Nomenclature

  • Chapter goals (three guiding questions):
    • What constitutes an atom?
    • How are elements classified in the periodic table?
    • How are binary and ionic compounds named?

Observing matter at the atomic scale

  • On the macroscopic scale we cannot see atoms with the naked eye; matter like a piece of copper appears uniform.
  • Powerful microscopy techniques allow visualization of matter at the atomic level:
    • Scanning Tunneling Microscopy (STM): uses tunneling data to draw pictures of surfaces. Example images show the surface of graphite (a form of elemental carbon) on the left and a sample of silicon on the right.
    • Atomic Force Microscopy (AFM): provides information about bonding between atoms; example image shows six-member carbon rings (benzene-like rings).
  • These tools enable us to infer atomic structure even though atoms are not visible optically.

Historical roots of the atomic idea

  • Democritus (≈400 BC): proposed matter is divisible only to a fundamental particle called atomos (indivisible).
  • Aristotle (ancillary contrast): proposed matter is composed of four elements (earth, fire, water, air); his influence delayed atomic theory for about 2,000 years.
  • Modern revival (early 1800s): John Dalton proposed a modern atomic theory with postulates; atoms reimagined as fundamental building blocks.

Dalton’s four postulates (early 1800s)

  • Postulate 1: Matter is composed of extremely small particles called atoms.
  • Postulate 2: All atoms of a given element are identical; atoms of different elements differ in fundamental ways.
  • Postulate 3: Atoms cannot be changed into atoms of another element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.
  • Postulate 4: Compounds are formed when atoms of more than one element combine; a given compound has a fixed ratio of constituent elements.
Laws that emerge from Dalton’s postulates
  • Law of constant (definite) composition: a pure compound always contains the same percentage by mass of each element, regardless of sample source or amount.
    • Example: calcium carbonate (CaCO₃, lime/limestone) is always about Ca ~ 40%, C ~ 12%, O ~ 48% by mass.
    • Example: water (H₂O) is always two hydrogens for every one oxygen.
  • Law of conservation of mass: mass is conserved in any process; atoms are rearranged but total mass remains the same.
    • Example: a reaction where two aqueous solutions (lead nitrate and sodium chromate) form a yellow solid (lead chromate) and a soluble product (sodium nitrate) still conserves mass overall.
  • Law of multiple proportions: if two elements form more than one compound, the masses of one element that combine with a fixed mass of the other element are in small whole-number ratios.

Illustrations of the laws (conservation and composition)

  • Law of constant composition illustrated with water and hydrogen peroxide:

    • Water:
    • Formula: $\mathrm{H_2O}$; atoms per molecule: 2 H, 1 O.
    • Mass ratio O:H = $\frac{16}{2} = 8:1$ (since O ≈ 16 amu and each H ≈ 1 amu; two hydrogens contribute 2 amu).
    • Hydrogen peroxide:
    • Formula: $\mathrm{H2O2}$; atoms per molecule: 2 H, 2 O.
    • Mass ratio O:H = $\frac{32}{2} = 16:1$ (oxygen mass 32 vs hydrogen mass 2).

  • Concept check: law of conservation of mass for the decomposition of hydrogen sulfide, $\mathrm{H2S\rightarrow H2 + S}$.

    • Given: starting mass of reactant = 6.5 g; mass of hydrogen produced = 0.384 g.
    • Mass of sulfur produced:
      m<em>S=m</em>reactants−m<em>H</em>2=6.5 extg−0.384 extg=6.116 g.m<em>{S} = m</em>{\text{reactants}} - m<em>{H</em>2} = 6.5\ ext{g} - 0.384\ ext{g} = 6.116\ \text{g}.

Development of modern atomic theory and subatomic particles

  • Dalton’s atomic model: atoms are hard, indivisible spheres (the "bowling ball" model).
  • Between 1895–1915 a series of experiments refined this picture:
    • Thomson’s plum pudding model (electrons embedded in a positively charged sphere).
    • Rutherford’s nuclear model: nucleus (protons and neutrons) at the center with electrons surrounding.
    • Bohr’s planetary model: electrons in specific orbits around the nucleus.
    • Schrödinger/Heisenberg/Einstein/Planck etc.: wave mechanics/quantum model; electrons occupy orbitals or electron clouds rather than fixed orbits.
  • Major subatomic particles:
    • Protons (p⁺) and neutrons (n⁰) in the nucleus; electrons (e⁻) orbiting the nucleus.
    • Protons: positive; neutrons: neutral; electrons: negative.
  • Subatomic discoveries and methods:
    • Cathode rays revealed electrons (Thomson’s work).
    • Radioactivity helped uncover subatomic particles.
    • Millikan oil-drop experiment refined electron charge/mass measurements.
    • Rutherford gold-foil experiment showed most of an atom’s mass is in the nucleus.
    • Bohr model (later refined by quantum theory) placed electrons in energy levels.
    • Chadwick discovered the neutron (1932).
  • Implications:
    • The nucleus holds most of an atom’s mass; electrons determine chemical behavior due to their involvement in chemical reactions.
    • Nuclear stability depends on neutron-to-proton ratio; too few or too many neutrons lead to instability and radioactivity.

Atomic structure and size relationships

  • The nucleus is the dense center; electrons form an surrounding cloud that defines the atom’s volume.
  • Relative mass and view:
    • The mass of the electron is much smaller than that of a proton or neutron, making the electron’s contribution to mass negligible in most calculations.
    • The nucleus accounts for most of the atom’s mass; the electron cloud accounts for the atom’s volume.
  • Spatial scales:
    • The diameter of a typical atom is about datom≈1.5 A˚.d_{\text{atom}} \approx 1.5\,\text{\AA}.
    • The diameter of the nucleus is about dnucleus≈10−4 A˚.d_{\text{nucleus}} \approx 10^{-4}\,\text{\AA}.
    • Therefore, the atom is roughly 10^4 times larger in diameter than its nucleus, illustrating that atoms are mostly empty space.
  • The Angstrom unit:
    • 1 A˚=1.0×10−10 m.1\ \text{\AA} = 1.0\times 10^{-10}\ \text{m}.
  • Masses:
    • Electron mass: me≈9.0×10−28 gm_e \approx 9.0\times 10^{-28}\ \text{g} (about 2000 times lighter than a proton or neutron).
    • Proton and neutron masses: approximately 1 amu each (the actual numbers are given in a table in your text as masses in amu).
    • Note: The electron’s small mass is what allows us to determine the charge/mass ratio and to deduce the masses of the other particles.
  • Quick dimensional-analysis check (optional practice from the slide):
    • If a bag weighs 1 lb, how many electrons does it contain?
    • Steps (conceptual):
    • Convert pounds to grams: $1\ \text{lb} \to \text{g}$.
    • Convert grams to kilograms as needed using $1\ \text{kg} = 10^3\ \text{g}$.
    • Use the electron mass in kilograms to find the number of electrons: N<em>e=m</em>bagme.N<em>e = \frac{m</em>{\text{bag}}}{m_e}.
    • Result (from the example): about 5×10295 \times 10^{29} electrons per pound.

The quantum view: atoms as more than just nuclei and planets

  • Atomic models history recap:
    • Dalton’s solid sphere (1803).
    • Thomson’s plum pudding (1897).
    • Rutherford’s nuclear model (early 1900s).
    • Bohr’s planetary model (early 1910s).
    • Schrödinger/Heisenberg/Planck/Pauli (1926 onward): quantum/ Wave-mechanical model.
  • The quantum model: electrons occupy orbitals or clouds rather than fixed orbits; the electron distribution defines chemical properties and behavior.

Symbols and neutral atoms; ions

  • The three subatomic particles and their symbols:
    • Electron: $e^-$, charge $-1$ (relative).
    • Proton: $p^+$, charge $+1$ (relative).
    • Neutron: $n^0$, charge $0$ (neutral).
  • Relative masses (in amu, approximate):
    • Protons and neutrons: about 1 amu each.
    • Electron: much smaller; often described as negligible in mass relative to the nucleus; the text notes $m_e \approx 9\times 10^{-28}$ g and that the electron is about 2000 times lighter than a proton or neutron.
  • Neutral atoms vs ions:
    • A neutral atom has an equal number of electrons and protons, giving no net electrical charge.
    • If the numbers are unbalanced, the species is an ion.

Putting it together: why this matters for chemistry

  • The nucleus determines the element’s identity (number of protons, $Z$).
  • The neutrons help stabilize the nucleus; too many or too few neutrons relative to protons can lead to instability and radioactivity.
  • The electrons determine chemical properties and reactivity through their arrangement and interactions in chemical bonds.
  • The atom is mostly empty space; its mass is concentrated in the nucleus, while the electron cloud defines the atom’s size and edge.

Quick reference of key constants and ideas (from the transcript)

  • Atomic units and sizes:
    • $d{\text{atom}} \approx 1.5\mathrm{\AA}$, $d{\text{nucleus}} \approx 10^{-4}\mathrm{\AA}$; ratio ~ $10^4$ in diameter.
    • $1\ \mathrm{\AA} = 1.0\times 10^{-10}\ \text{m}$.
  • Subatomic particle masses (as stated in the notes):
    • Electron mass: $m_e \approx 9\times 10^{-28}\ \text{g}$.
    • Protons and neutrons: mass ≈ 1 amu each (as described in the summary table).
  • Common symbols and charges:
    • $e^-$ (−), $p^+$ (+), $n^0$ (0).

Note: The text also points toward later chapters for deeper explorations of atomic models and the full quantum mechanical treatment of electrons in atoms. The current notes summarize the material up through the introduction to atomic theory and the basic structure of atoms, preparing you for more detailed studies in later chapters.