A t o m i c P h y s i c s What is the ATOM???

1. AtomicPhysicsWhat is the ATOM???

2. MATTER=ATOM • All matter is composed of atoms. • Atoms are the smallest part of an element that keeps that element’s properties. • If an atom were the size of a football field, the nucleus would be the size of a marble and nearly all of the mass of the atom is in the nucleus. • The space around the nucleus and the electron have very, very little mass. This means that matter is mostly empty!

3. The Atom and the nucleus The nucleus contains most of the mass of the atom (protons + neutrons). The rest is mostly empty space!

4. DemocritusThe Greek Understanding • Democritus (BC 400) and the “Atomists” debate the “four” elements (fire, water, earth, and air). • Democritus concluded that matter could not be divided into smaller and smaller pieces forever. Eventually, the smallest piece of matter would be found.

5. Democritus (400 B.C.) • He used the word “atomos” to describe the smallest possible piece of matter. • Not based on experimental data

6. Alchemy (next 2000 years) • Mixture of science and mysticism. • Lab procedures were developed, but alchemists did not perform controlled experiments like true scientists.

7. John Dalton • John Dalton (1766-1844) proposed an atomic theory • British Schoolteacher • based his theory on others’ experimental data • Billiard Ball Model • atom is a uniform, solid sphere

8. Dalton’s Four Postulates • All elements are composed of indivisible particles called atoms. • Atoms of the same element are exactly alike (isotopes not known yet). • Atoms of different elements are different. • Compounds are formed by joining atoms of two or more elements.   • In a chemical reaction, atoms are rearranged, but not changed.

9. Dalton’s Atomic Theory While this theory was not completely correct, it revolutionized how chemists looked at matter and brought about chemistry as we know it today instead of alchemy Thus, it’s an important landmark in the history of science. Law of definity proportions law of multiple proportions

10. Problems with Dalton’s Atomic Theory? 1. matter is composed, indivisible particles Atoms Can Be Divided, but only in a nuclear reaction 2. all atoms of a particular element are identical Does Not Account for Isotopes (atoms of the same element but a different mass due to a different number of neutrons)! 3. different elements have different atoms YES! 4. atoms combine in certain whole-number ratios YES! Called the Law of Definite Proportions 5. In a chemical reaction, atoms are merely rearranged to form new compounds; they are not created, destroyed, or changed into atoms of any other elements. Yes, except for nuclear reactions that can change atoms of one element to a different element

11. Henri Becquerel (1896) • Discovered radioactivity • spontaneous emission of radiation from the nucleus • Three types: • alpha () - positive • beta () - negative • gamma () - neutral

12. Thompson’s “Plum Pudding” Model • Discovered negatively charged particles coming off a gas known to be neutral. • Called these “early” electrons, corpuscles. • Predicted that there must be positive particles (protons) present in equal number to balance negative charges. • Proposed atoms were made of pudding like positive charges, with negative electrons inside scattered like plums.

13. J. J. Thomson (1903) • Cathode Ray Tube Experiments • beam of negative particles • Discovered Electrons • negative particles within the atom • Plum-pudding Model

16. J. J. Thomson (1903) Plum-pudding Model • positive sphere (pudding) with negative electrons (plums) dispersed throughout

17. Ernest Rutherford (1911) • Gold Foil Experiment • Discovered the nucleus • dense, positive charge in the center of the atom • Nuclear Model • 99.9% of total mass is in nucleus

18. Rutherford’s Model • Fired tiny positive particles at gold foil…most went through without change of course. What did this prove? • Some particles bounced back as if they hit something solid. • Reasoned that they were repelled by similar charged particles. • Discovered and proved the existence of the proton. • Reasoned that atoms are mostly “empty” space, but with a small dense center with positive charges.

19. Rutherford’s experiment.

20. Rutherford’s Gold Foil Experiment

21. Results of foil experiment if Plum Pudding model had been correct.

22. What Actually Happened

23. Ernest Rutherford (1911) • Nuclear Model • dense, positive nucleus surrounded by negative electrons

24. These deflections were not consistent with Thomson's model. Rutherford was forced to discard the Plum Pudding model. • He reasoned that the only way the alpha particles (positively charged) could be deflected backwards was if most of the mass in an atom was concentrated in a nucleus.

25. Robert A. Millikan • Experiments showed that the mass of an electron is smaller than the simplest type of hydrogen atom. • Mass of electron = 9.109 x 10-31kg • Electron has a negative charge.

26. Based on what was learned about electrons, • Two other inferences were made about atomic structure: • 1. Because atoms are electrically neutral, they must contain a positive charge to balance the negative electrons. • 2. Because electrons have so much less mass than atoms, atoms must contain other particles that account for most of their mass.

27. Mosely • Determined the number of PROTONS in elements

28. The Bohr Model of the Atom • Neils Bohr suggested that only certain electron orbits are allowed. • Orbits are given number n=1,2,3,4,5,…infinity. • n=1 is the ground state (ground orbit) and has the lowest energy. No lower orbit is permitted. • The excited states (higher orbits) have more energy. • The orbits get closer together as n increases. • To change orbits, the electron must either gain or loose energy. • Loss of energy appears as light • What we see as light is actually the release of a photon • Each type of atom has a unique set of energy levels.

29. Niels Bohr (1913) • Bright-Line Spectrum • tried to explain presence of specific colors in hydrogen’s spectrum • Energy Levels • electrons can only exist in specific energy states • Planetary Model • Based on Hydrogen

30. Niels Bohr (1913) • Planetary Model • electrons move in circular orbits within specific energy levels Bright-line spectrum

31. Energy of photon depends on the difference in energy levels Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom Bohr Model 6 5 4 3 2 1

32. Neils Bohr’s Model

33. Erwin Schrodinger (1926) • Quantum mechanics • electrons can only exist in specified energy states • Electron cloud model • orbital: region around the nucleus where e- are likely to be found

34. Erwin Schrodinger’s Model • Developed wave Quantum Mechanical Model. • Treats atoms as a 3 dimensional system of waves. • Contains ideas of Bohr’s and incorporates others.

35. Quantum Mechanical Model • Quantum numbers specify the electrons probable location and energy level • Because electrons are always moving their exact whereabouts are unknown

36. Erwin Schrödinger (1926) Electron Cloud Model (orbital) • dots represent probability of finding an e-not actual electrons

37. Erwin Schrodinger’s Model • Developed wave Quantum Mechanical Model. • Treats atoms as a 3 dimensional system of waves. • Contains ideas of Bohr’s and incorporates others.

38. James Chadwick (1932) • Discovered neutrons • neutral particles in the nucleus of an atom • Joliot-Curie Experiments • based his theory on their experimental evidence

39. James Chadwick (1932) Neutron Model • revision of Rutherford’s Nuclear Model