CHEM101Ch2

Electrons and Early Atomic Theories

J. J. Thomson (1897)

• Discovered electrons using cathode ray tubes, demonstrating that the rays produced were composed of negatively charged particles.

• He found that cathode rays were deflected toward positively charged plates, leading to the conclusion that these particles are negatively charged.

• Established that atoms contain negatively charged particles with a constant mass-to-charge ratio of approximately 1.76 × 10^8 C/kg.

Millikan’s Oil Drop Experiment

• Conducted by Robert Millikan in 1909 to determine the charge and mass of electrons.

• The charge of an electron (e) was measured to be e = −1.602 × 10−19 C.

• The mass of an electron (me) was determined to be approximately 9.109 × 10−31 kg through the measurement of the droplets' movement in an electric field.

Plum-Pudding Model

• Proposed by Thomson, it suggested that atoms consist of a diffuse, positively charged sphere within which negatively charged electrons are embedded, akin to raisins in a pudding.

Radioactivity

• Discovered by Henri Becquerel in 1896 when he found that pitchblende emitted radiation.

• The work was extended by Marie and Pierre Curie, and Ernest Rutherford, identifying three main types of radiation:

  • Beta particles (β): high-energy, high-speed electrons emitted from certain types of radioactive nuclei.

  • Alpha particles (α): positively charged particles (equivalent to the helium nucleus, 2+ charge).

Rutherford’s Gold Foil Experiment

• Involved bombarding a thin gold foil with alpha particles to test the validity of Thomson's model of the atom.

• Results indicated a significant majority of particles passed through, but some were deflected at large angles, leading to the conclusion that

  • Atoms have a small, dense nucleus that is positively charged, containing most of the atom's mass.

  • Identified electrons existing within a diffuse electron cloud surrounding the nucleus.

Concept of the Nuclear Atom

• Characterized by a tiny, positively charged nucleus that holds protons and neutrons, with electrons occupying the surrounding space.

• Proton: Subatomic particle with a positive charge and mass approximately 1.6726 × 10−27 kg.

• Neutron: Subatomic particle without charge, mass approximately 1.6750 × 10−27 kg.

Subatomic Particles and Atomic Mass

• Masses of atoms and subatomic particles are measured in unified atomic mass units (u), where 1 u = 1.66053906660 × 10−27 kg.

• Carbon is often used as a reference with a standard atomic mass of 12 u (6 protons + 6 neutrons).

Isotopes

• Definition: Atoms of the same element possessing the same number of protons but differing numbers of neutrons, leading to variations in atomic mass.

• Evidence from Aston’s experiment utilized a mass spectrometer to reveal multiple spots in a neon spectrum, confirming the existence of isotopes.

Symbol Representation of Isotopes or Nuclides

• Nuclide Symbol: A = mass number (total number of nucleons), Z = atomic number (number of protons), X = symbol of the element.

• Example of Isotopes:

  • 22 10Ne: (10 protons, 12 neutrons).

  • 20 10Ne: (10 protons, 10 neutrons).

Exercises and Applications

• Exercises focus on relating symbols of nuclides to their nuclei composition and identifying atomic structure based on given numbers of protons and neutrons.

Mendeleev’s and Modern Periodic Table

• Mendeleev's Table was initially arranged by increasing atomic mass, while leaving gaps for undiscovered elements.

• The Modern Table organizes elements by increasing atomic number, reflecting advancements in the understanding of atomic structure.

Element Categories

• Metals: Typically shiny, malleable, good conductors of electricity (notable exception: mercury - liquid at room temperature).

• Nonmetals: Vary considerably in state at room temperature and are generally poor conductors.

• Metalloids: Display metalloid properties, being brittle yet shiny and acting as semiconductors.

Common Group Names

• Group 1: Alkali metals (e.g., Lithium, Sodium).

• Group 2: Alkaline earth metals (e.g., Magnesium, Calcium).

• Group 16: Chalcogens (e.g., Oxygen, Sulfur).

• Group 17: Halogens (e.g., Fluorine, Chlorine).

• Group 18: Noble gases (e.g., Helium, Neon).

Molar Mass and Avogadro's Constant

• Molar Mass: Mass of a substance per mole, numerically equivalent to the atomic mass in unified atomic mass units.

• Avogadro's Constant: 1 mole contains approximately 6.022 × 10^23 particles, a fundamental constant in chemistry.

Mass Spectrometry

• The process involves converting atoms into ions (M+) which are then sorted by their mass-to-charge ratio.

• The resulting mass spectrum displays intensity versus m/z (mass-to-charge) ratios, allowing for the identification of various species.

• Example: The explosive compound TATP can be analyzed through its molecular-ion peak observed at a specific m/z value in the mass spectrum.