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THE ATOM

THE ATOM. Objectives: Understand the experimental design and conclusions used in the development of modern atomic theory, including Dalton’s Postulates, Thomson’s discovery of electron properties, Rutherford’s nuclear atom, and Bohr’s nuclear atom. Democritus.

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THE ATOM

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  1. THE ATOM Objectives: Understand the experimental design and conclusions used in the development of modern atomic theory, including Dalton’s Postulates, Thomson’s discovery of electron properties, Rutherford’s nuclear atom, and Bohr’s nuclear atom.

  2. Democritus • Made his discovery around the year 250 B. C. • This was the first discovery about the atom, the next would come in another 2000 years.

  3. The First Atom • Democritus took a sea shell and broke it in half. • Than he broke it in half again. • When the pieces got to small he use a mortar and pestle to crush the shell. • He finally believed he got to the smallest piece possible and called it the ATOM; which in Greek means INDIVISIBLE.

  4. John Dalton (1766-1844)

  5. A New System of Chemical Philosophy (1808)

  6. Dalton’s Atom Model • All matter is made on atoms; and atoms are indivisible. • Atoms of the same element are all identical. • Compounds are formed by a combination of two or more different atoms and they always have the same proportion of elements. THE LAW OF DEFINITE COMPOSITION • A chemical reaction is a rearrangement of atoms and the atoms are neither created nor destroyed. THE LAW OF CONSERVATION OF MATTER

  7. J. J. Thomson (1856-1940)Joseph John Thomson • English physicist who in 1897 discovered a particle smaller than the atom ; the electron. • Particle has a negative charge and is much smaller than the atom so must come from the inside of the atom. • Electrons are scattered around the atom like raisins in pudding. (THE PLUM PUDDING MODEL)

  8. Thomson and Rutherford

  9. Rutherford’s GoldFoil Experiment

  10. Rutherford’s GoldFoil Experiment

  11. New Zealand born physicist; worked in England 1911 conducted the “Gold Foil Experiment” the proved the existence of a small positively charged center of the atom. Disproved the “Plum Pudding Model” THE NUCLEAR MODEL Discovered the proton. Thought that the electrons orbited the nucleus like planets orbited the sun. Ernest Rutherford (1871-1937)

  12. Millikan’s Oil Drop Experiment • A fine mist of oil droplets is introduced into the chamber. • The oil is ionized by x-rays. • The electrons adhere to the oil drops. • The value for the charge of the electron can be calculated.

  13. Niels Bohr (1885-1962) • Danish physicist, produced his model in 1911. • Saw problems with Rutherford’s model. • If electrons “orbit” than they are changing direction so they are accelerating. • That would require energy.

  14. The Orbital Model • Electrons do not “orbit” but are in allowable ENERGY LEVELS. • When the electrons stay in these levels, which are at specific distances from the nucleus, they do not give off energy.

  15. Bright Line Spectrum • But, if the electron moves from one level to another it gives off or absorbs energy. • These Bright Line Spectrums are produced when the electrons “fall back” to a lower energy level and give off energy. • Every element has a unique Bright Line Spectrum.

  16. Subatomic Particles Objective: Be able to calculate the number of protons, neutrons, and electrons in any atom, ion, or isotope.

  17. The Subatomic Particles THE PROTON • p+ • positively charged • located in the nucleus • relative mass = 1 atomic mass unit • mass = 1.673 x 10-24 grams • equal to atomic number • number of protons “defines” the atom

  18. The Subatomic ParticlesTHE NEUTRON • n0 • neutral (no electrical) charge • located in the nucleus • relative mass = 1 atomic mass unit • mass = 1.675 x 10-24 grams • equal to mass number minus atomic number • James Chadwick proposed the existence of the neutron.

  19. The Subatomic ParticlesTHE NEUTRON • Isotopes – different atoms of the same element that have the same number of protons but different numbers of neutrons • some isotopes are radioactive – they emit energy when the nucleus of the atom breaks down spontaneously • most radioactive isotopes are not dangerous • to determine if an isotope is radioactive calculate the proton to neutron ratio • if ratio is greater than or less than 1:1 for “small” atoms the isotope is unstable (smaller than Ca) • if ratio is greater then 1:1.5 for “large” atoms the isotope is unstable

  20. The Subatomic ParticlesTHE ELECTRON • e- • located in the electron cloud which is divided into energy levels, sublevels, orbitals, and spins • relative mass = 0 atomic mass units • mass = 9.11 x 10-28 grams • equal to the number of protons if atom is neutral • atom becomes a charged ion if electrons are gained or lost • positive ion = CATION • formed by the loss of electron, happens to metals • negative ion = ANION • formed by the gain of electron, happens to nonmetals

  21. Location of Electrons • Energy Levels • Discovered by Niels Bohr • # electrons = 2n2 • “n” is the energy level • 1st level can hold 2 e- • 2nd level can hold 8 e- • 3rd level can hold 18 e- • (eight if the outside energy level) • 4th level can hold 32 e- • (eight if the outside) • The outside level is called the valance level and can never hold more than 8 electrons.

  22. NUCLEAR SYMBOLS mass number ion charge 23 +1 Na p+ = 11 11 n0 = 12 atomic number e- = 10

  23. Objective • Use isotopic composition to calculate the average atomic mass of an element.

  24. Mass Number vs. Atomic Mass • mass is given for individual atoms • mass number is given in nuclear symbols • atomic mass is an average mass for all isotopes for the element • atomic mass is the number on the periodic table • if you round the average atomic mass you will have the mass number of the most common isotope

  25. Average Atomic Mass

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