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 • 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) • 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 • 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 • 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 • 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 • 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 • 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.
Radioactivity • 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 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 • 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 • 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.