Discovery of Atomic Structure

1. Discovery of Atomic Structure • By 1850 scientists knew that atoms were composed of charged particles. • Electrostatic attraction: • Like charges repel • Opposites attract Mullis

2. Cathode Rays and Electrons • C.R. 1st discovered in mid-1980s from studies of electrical discharge thru partially evacuated tubes (CRTs) • Cathode rays = radiation produced when high voltage is applied across the tube. • The voltage causes negative particles to move from the negative electrode (cathode) to the positive electrode (anode). • The path of electrons can be altered by the presence of a magnetic field. Mullis

3. Consider cathode rays leaving the positive electrode through a small hole…. • If they interact with a magnetic field perpendicular to an applied electric field, then the cathode rays can be deflected by different amounts. • Amount of deflection depends on applied magnetic and electric fields. • Deflection also depends on the charge-to-mass ratio of an electron. • Thomson determined the charge-to-mass ratio of an electron in 1897. • Charge-to-mass ratio = 1.76 x 108 C/g • C: Coulomb, SI unit of electric charge Mullis

4. Millikan Oil-Drop Experiment • Sprayed oil drops over the hole in a positively charged plate and measured the electrostatic force of attraction. • Found the charge on the electron to determine its mass • Concluded the charge on the electron must be 1.60 x 10-19 C • Mass of electron = 1.60 x 10-19 C = 9.10 x 10 -28 g 1.76 x 108 C/g Mullis

5. Radioactivity(Spontaneous emission of radiation) Mullis

6. A Positively Charged Nucleus • Rutherford shot alpha particles though a thin piece of gold foil. • Some of these particles were deflected instead of passing straight through • Recall “like repels like.” • When a + alpha particle encountered a nucleus of a gold atom, it was deflected by the dense positively charged nucleus. Mullis

7. Scientist Contributions • Thomson: • “Discovered” electron (1897) • Cathode ray experiments • “Plum pudding” atomic model • Millikan: • Mass of electron • Oil-drop experiment (1909) • Rutherford: • Positively charged nucleus (1911) • Gold foil experiments • Discovered proton (1919) • Chadwick: Discovered neutron (1932) Mullis

8. Small Numbers • Electronic Charge: 1.609 x 10-19 C • Charge on an electron: -1.609 x 10-19 C • Charge on a proton: +1.609 x 10-19 C • Atomic Mass Unit (amu): 1.66054 x 10-24 g • Proton mass: 1.0073 amu • Neutron mass: 1.0087 amu • Electron mass: 5.486 x 10-4 amu • Unit of length used to note atomic dimensions: 1 Angstrom(Å) = 1x10-10 m Mullis

9. Atomic Number • Number of protons or electrons in an element • Identifies the element • Atomic Mass • Nucleus contains most of the mass of an atom. • Protons and neutrons are each ~ 1.67 x 10-24 g. • Electrons are each ~ 9.11 x 10-28 g. • Use atomic mass unit (amu) instead of gram. • The mass of one proton is ~ 1 amu. • Mass Number • The sum of the number of protons and number of neutrons in the nucleus • Is approximately equal to the average atomic mass shown on periodic table. • Number of neutrons = mass number – atomic number Mullis

10. Isotopes • Atoms of the same element with different numbers of neutrons • Have the same number of protons • Example: Carbon-12 and Carbon-14 • Radioactive Isotopes • Unstable in nature • Can be used to date fossils and rocks • The time it takes for half of the radioactive atoms in a piece of the fossil to change to another element is its half-life. Mullis

11. Isotopes AX Z • Isotopes have the same Z, but different A. • Isotopes have different numbers of neutrons. • An atom of a specific isotope is called a nuclide. • Nuclidesof hydrogen include: • 1H = hydrogen (protium) • 2H = deuterium (heavy hydrogen) • 3H = tritium (3H is radioactive.) Mullis