The Structure of Atoms: Discovering Electrons, Protons, and Neutrons

1. CHAPTER 5 • The Structure of Atoms

2. Fundamental Particles • Three fundamental particles make up atoms:

3. The Discovery of Electrons • Late 1800’s & early 1900’s Cathode ray tube experiments showed that very small negatively charged particles are emitted by the cathode material. • 1897 – J. J. Thomson Modified the cathode ray tube and measured the charge to mass ratio of these particles. He called them electrons. (Nobel prize in physics, 1906)

4. The Discovery of Electrons • 1909 – Robert A. Millikan Determined the charge and the mass of the electron from the oil drop experiment. (The second American to win Nobel prize in physics in 1923) • 1910 – Ernest Rutherford Gave the first basically correct picture of the atom’s structure. (Nobel prize in chemistry in 1908)

5. Rutherford’s Atom • The atom is mostly empty space • It contains a very small, dense center called the nucleus • Nearly all of the atom’s mass is in the nucleus • The nuclear diameter is 1/10,000 to 1/100,000 times less than atom’s radius

6. The Discovery of Protons • 1913 – H.G.J. Moseley Realized that the atomic number defines the element: • Each element differs from the preceding element by having one more positive charge in its nucleus • Along with a number of observations made by Rutherford and some other physicists, this led to the discovery of the proton • The elements differ from each other by the number of protons in the nucleus

7. The Discovery of Neutrons • 1932 – James Chadwick recognized existence of massive neutral particles which he called neutrons (Nobel prize in physics in 1935) • The atomic mass of an element is mainly determined by the total number of protons and neutrons in the nucleus • The atomic number of an element is determined by the total number of protons in the nucleus

8. Mass Number and Atomic Number • Mass number – A • Atomic number – Z • Z = # protons • A = # protons + # neutrons • # protons = # electrons • The way we typically write this: full nuclide symbol short nuclide symbol

9. Isotopes • Atoms of the same element but with different masses • The same element means that the number of protons is the same, • then different masses mean that the number of neutrons differs protium (or hydrogen) deuterium tritium

10. Isotopes: Example

11. Experimental Detection of Isotopes • 1919-1920 – Francis Aston Designed the first mass-spectrometer (Nobel prize in chemistry in 1922) • Factors which determine a particle’s path in the mass spectrometer: • accelerating voltage, V • magnetic field strength, H • mass of the particle, m • charge on the particle, q

12. Mass Spectrometry & Isotopes • Mass spectrum of Ne+ ions • This is how scientists determine the masses and abundances of the isotopes of an element

13. Mass Spectrometry & Isotopes • Let’s calculate the atomic mass of Ne using the mass-spectrometry data

14. Atomic Weight Scale • A unit of atomic mass (atomic mass unit) was defined as exactly 1/12 of the mass of a 12C atom • Two important consequences of such scale choice: • The atomic mass of 12C equals 12 a.m.u. • 1 a.m.u. is approximately the mass of one atom of 1H, the lightest isotope of the element with the lowest mass. • The atomic weight of an element is the weighted average of the masses of its isotopes

15. Isotopes and Atomic Weight • Naturally occurring chromium consists of four isotopes. It is 4.31% 50Cr, mass = 49.946 amu 83.76% 52Cr, mass = 51.941 amu 9.55% 53Cr, mass = 52.941 amu 2.38% 54Cr, mass = 53.939 amu Calculate the atomic weight of chromium

16. Isotopes and Atomic Weight • Naturally occurring Cu consists of 2 isotopes. It is 69.1% 63Cu with a mass of 62.9 amu, and 30.9% 65Cu, which has a mass of 64.9 amu. Calculate the atomic weight of Cu to one decimal place. • A.W.(Cu) = (62.9 amu  0.691) + (64.9 amu  0.309) = = 63.5 amu

17. Electromagnetic Radiation • Any wave is characterized by 2 parameters: • Wavelength () is the distance between two identical points of adjacent waves, for example between their crests It is measured in units of distance (m, cm, Å) • Frequency () is the number of wave crests passing a given point per unit time (for example, per second) It is measured in units of 1/time, usually s-1 1 s-1 = 1 Hz (Hertz)

18. Electromagnetic Radiation • The speed at which the wave propagates: c =  • The speed of electromagnetic waves in vacuum has a constant value: c = 3.00108 m/s • This is the speed of light • Given the frequency of the electromagnetic radiation, we can calculate its wavelength, and vice versa

19. Electromagnetic Radiation • Max Planck (Nobel prize in physics in 1918) • Electromagnetic radiation can also be described in terms of “particles” called photons • Each photon is a particular amount of energy carried by the wave • Planck’s equation relates the energy of the photon to the frequency of radiation: E = h  (h is a Planck’s constant, 6.626·10-34 J·s)

20. Electromagnetic Radiation • What is the energy of green light of wavelength 5200 Å?