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Structure of the Atom. Lecture 1. Early Theories of Matter. Democritus (460-370 B.C.) – Greek Philosopher Proposed that matter was not infinitely divisible Matter was made up of tiny individual particles called atomos Atomos could not be created, destroyed, or further divided

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early theories of matter
Early Theories of Matter
  • Democritus (460-370 B.C.) – Greek Philosopher
    • Proposed that matter was not infinitely divisible
    • Matter was made up of tiny individual particles called atomos
    • Atomos could not be created, destroyed, or further divided
    • Different kinds of atomos had different sizes and shapes
    • Changes in matter resulted from the change in the groupings of atomos and not from the atomos themselves.

Aristotle (384-322 B.C.) – Greek Philosopher

    • Rejected Democritus’ atomic theory because it did not agree with his own ideas on nature
    • He did not believe that the “nothingness” of empty space through which atomos moved could exist
john dalton 1766 1844
John Dalton (1766-1844)
  • Revived and revised Democritus’ ideas based upon his own scientific research.
  • Devised his own atomic theory which included
    • All matter is composed of extremely small particles called atoms
    • All atoms of a given element are identical, having the same size, mass, and chemical properties
    • Atoms of a specific element are different from those of any other element
    • Atoms cannot be created, divided into smaller particles, or destroyed
    • Different atoms combine in simple whole-number ratios to form compounds
    • In a chemical reaction, atoms are separated, combined, or rearranged.
atoms defined
Atoms: Defined
  • An atom is the smallest particle of an element that retains the properties of the element.
  • Atoms are composed of subatomic particles found in the nucleus of the atom and outside the nucleus.
    • Protons: positively charged, mass of 1 amu, found in the nucleus
    • Neutrons, no charge, mas of 1 amu, found in the nucleus
    • Electrons: negatively charged, mass of 1/1840thamu, found outside the nucleus in electron valences
discovering the electron
Discovering the Electron
  • J.J. Thomson (1890s) conducted cathode ray (streams of negatively charged particles) experiments to determine the ratio of the charged cathode ray particles to their mass.
    • He concluded that the mass of the particle was less than that of a hydrogen atom.
    • He proposed that the atom was composed of a uniform positive charge that had negatively charged electrons distributed throughout. It was called the Plum-pudding model.
  • Robert Millikan (1909) determined the mass of an electron to be 9.1 x 10-28 g or 1/1840th the mass of a hydrogen atom.
nuclear atom
Nuclear Atom
  • Ernest Rutherford (1911) conducted an experiment to determine how positively charge alpha particles interacted with solid matter.
    • Alpha particles have the mass of a Helium atom and contain 2 protons and 2 neutrons, but no electrons.
  • He suspected only minor deflections of the particles based on the Plum-pudding model. He believed that the alpha particles would shoot through matter without interacting with the matter’s atoms.
  • The alpha particles were sharply deflected back toward the source of the alpha particles or at sharp angles away from the matter.

Rutherford concluded that the atom was composed of a tiny, dense region, he called the nucleus, which contained all of an atom’s positive charge and mass surrounded by mostly empty space containing the electrons.

  • He hypothesized that electrons were held in place orbiting the nucleus by their attraction to its positive charge.
  • By 1920, Rutherford concluded that the nucleus contained positively charge particles he called protons.
  • Protons have a mas of 1.673 x 10-24 g or 1 amu.

James Chadwick (1932) showed that the nucleus also contained another subatomic particle, a neutral particle he called the neutron.

    • A neutron has a mass nearly equal to that of a proton (1.675x10-24 g or 1 amu).
atomic number
Atomic Number
  • Henry Mosely (1887-1915) discovered that each element contained a unique positive charge in its nucleus.
  • The number of protons identifies each particular element.
  • The number of protons is the element’s atomic number.
  • All atoms are neutral and the number of protons must equal the number of electrons.
isotopes and mass number
Isotopes and Mass Number
  • Atoms with differing numbers of neutrons are called isotopes.
    • Isotopes have the same number of protons and eletrons, just differing numbers of neutrons.
  • The mass number of an element is the AVERAGE of the sum of the protons and neutrons in the nucleus of all the isotopes of that element.
  • Isotopes are given the elemental name and the mass number for that specific isotope.
    • For example: potassium with 19 protons and 20 neutrons would be Potassium-39 or K-39.
  • Changes that involve the nucleus of an atom are referred to as nuclear reactions.
  • Some atoms spontaneously emit radiation in a process called radioactivity.
  • Radioactive atoms undergo significant changes that can alter their identities.
    • By emitting radiation, atoms of one element transform into atoms of a different element.
radioactive decay
Radioactive Decay
  • Radioactive elements emit radiation because their nuclei are unstable.
  • This process is called radioactive decay.
  • Atoms undergo radioactive decay until they become stable non-radioactive atoms, resulting in a different element.
types of radiation
Types of Radiation
  • Alpha: attracted toward negatively charged particles
    • Contain two protons and two neutrons
    • Have a 2+ charge and a mass of 4 amu
    • Is equivalent to a helium nucleus
  • Beta: attracted toward positively charged particles
    • Are fast moving electrons
    • Beta decay of Carbon-14 results in Nitrogen and 1 beta particle
  • Gamma radiation: high energy rays that have no mass and no charge
    • Accompany alpha and beta particle emission and account for most of the energy lost during radioactive decay.
nuclear stability
Nuclear Stability
  • The ratio of neutrons to protons is a primary indicator of an atom’s stability.
    • Atoms containing too many or too few neutrons are unstable.
    • Alpha and beta emissions affect the neutron to proton ratio of a newly created nucleus.
    • There are few radioactive atoms in nature because most of them have already decayed into stable atoms.