ATOMIC THEORIES

1. ~ ~ Fire WaterEarth Air ATOMIC THEORIES SCH4U Date:_______________________

2. Democritus ( ~300 BC) • Hypothesized that matter cur into smaller and smaller pieces would eventually reach the ‘atomos’, meaning indivisible. • Atoms were small, hard particles of the same material, but were different shapes and sizes • Theory was ignored and forgotten for about 2000 years.

3. GOLD SILVER COPPER IRON SAND . . . . . . . . . . . . . . . Alchemy(500 – 1400 A.D.) Alchemical symbols for substances… Transmutation: changing one substance into another In ordinary chemistry, we cannot transmute elements.

4. Contributions of Alchemists • Information about elements The elements mercury, sulfur and antimony were discovered Propoerties of some elements • Develop lab apparatus/procedures/ experimental techniques- learned how to prepare acids - developed several alloys-new glassward

5. Aristotle and Plato’s Theories • 4 elements, replaced the theory of the atom until the 1800’s

6. John Dalton (1805) • Recreated the modern theory of atoms to explain three important scientific laws.

7. Dalton’s Theory Cont’d • Each element is composed of extremely small particles called atoms • All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements. • The atoms of one element cannot be changed into atoms of a different element by chemical reactions; atoms are neither created or destroyed • Compounds are formed when atoms of more than one element combine, a given compound always has the same relative number of atoms.

8. J.J Thompson (1897) • Known for the Plum Pudding Model or Raisin-Bun Model - Dough was positive, Raisins were negative Positive Sphere with embedded electrons.

9. Used Cathode-Ray Tubes to study the passage of an electric current through a gas. As the current passed, it gave off rays of negatively charged particles. • Research group at Cambridge University in England used mathematics to predict the uniform three-dimensional distribution of electrons (corpuscles) throughout an atom. • Robert Millikan later measured and determined the charge on an electron.

10. Charge & Mass of an electron • 1.60217662 × 10-19coulombs Electron Mass = 1.602 x 10 -19 C = 9.10 x 10-28 g 1.76 x 108 C/g Thomsons Charge-to-Mass Ration 1.76 x 108 C/g

11. Ernest Rutherford • J.J Thomson’s student • Studied nuclear radiation for nine years at McGill University in Montreal - alpha(helium nuclei) - beta - gamma • Received a Nobel Prize in Chemistry for his work with radioactivity

12. Gold-Foil Experiment • Used Ra as a source of alpha radiation which was directed at thin film of gold (~2000 atoms thick). • Prediction: based on Thomson model that alpha particle should not be deflected (or very little)

13. Rutherford’s Apparatus beam of alpha particles radioactive substance circular ZnS - coated fluorescent screen gold foil Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120

14. Results of foil experiment if plum-pudding had been correct. Electrons scattered throughout positive charges - + + - + + - + - - + + - + - - Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 57

15. What he expected…

16. What he got… richocheting alpha particles

17. Rutherford’sGold Foil Experiment (1909) Revised Theory

18. Interpreting the Observed Deflections • Atom is mostly empty space • Small dense, positive piece at center (nucleus) • Electrons float around the nucleus • Electrostatic force is directly proportional to the quantity of electric charge involved. A greater charge exerts a greater force. The force exerted on an alpha particle by a concentrated nucleus would be much greater that the force exerted on an alpha particle by a single proton. Hence, larger deflections will result from a dense nucleus than from an atom with diffuse positive charges.

19. Rutherford used physics to calculate how small the nucleus would have to be produce the large-angle deflections observed. He calculated that the maximum possible size of the nucleus is about 1/10,000 the diameter of the atom. Rutherford concluded that the atom is mostly space • Concluded that a nuclear force of attraction has to much stronger than the electrostatic force repelling the positive charges in the nucleus • Later, work of Rutherford and Chadwick resulted in determining the existence of a positive particle called a proton • In 1932, Chadwick determined the existence of a neutral particle – the neutron.

20. Planck’s Quantum Hypothesis • Credited with starting the quantum revolution by surprising interpretation of study of light. • His teacher, Gustav Kirchoff – studied light emitted from “black bodies”. • A blackbody is used to describe an idea, perfectly black object that does not reflect any light, and emits various forms of light as a result of its temperature.

21. Blackbody Effect • As a solid is heated to high temperaure, it begins to glow ( red  white) . Recall that white light is a combination of all colours- so light emitted by the hotter object should emit all coloured light.

22. When electronic instruments are used to measure intensity (brightness) of the different colours observed in the spectrum that light emitted, a typical bell shaped curve is observed Notice the curve becomes higher and shifts towards the higher-energy UV as the temperature increases

23. Plank hypothesized that energies of the oscillating atoms in the heated solid were multiples of a small quantity of energy, and not continuous. • Plank was reluctant to purse this line of reasoning, Albert Einstein later pointed out that the inevitable conclusion of Planck’s hypothesis that light emitted by a hot object is also quantized. • One little burst of packet of energy is known as a quantum.

24. The smallest amount of energy, a quantum, is given by: E = hv, where h is Planck’s constant: = 6.626 × 10–34 J s Therefore, as temperature is increased, the proportion of each larger quantum becomes greater. The colour of a heated object is due to a complex combination of the number and kind of quanta.

25. Rise of the Quantum Theory • Green philosophers believed that light was a stream of particles • In the late 17thcentury, Christian Huygens proposed that light can be best explained as a wave (opposed by Isaac Newton). • Observations of refraction, reflection and diffraction favoured the wave hypothesis

26. Classical Theory of Light • Mid-19th century, James Maxwell proposed a brilliant theory, explaining light as an electromagnetic waved, composed of electric and magnetic fields that can exert forces.

27. Characterizing Waves • Frequency of a wave is the number of cycles of the wave that pass through a point in a unit of time • Amplitude of a wave is its height: the distance from a line of no disturbance through the center of the wave peak

28. The Photoelectric Effect • Heinrich Hertz discovered the photoelectric effect by accident in 1887 • It is the release of electrons from a substance due to light striking the surface of a metal. • According to classical theory, the brightness (intensity) of the light shone on the metal, would determine the kinetic energy of the liberated atoms. The brighter the light the greater the energy of the liberated electrons. FALSE !!!

29. The Photoelectric Effect • The frequency(colour/energy) of the light was the most important characteristic in producing the effect. • Albert Einstein (1906) used Planck’s idea of quantum to reason that light consisted of a stream of energy packets or quanta- called photons. • The energy of the photon is transferred to the electron. Some of this energy is used by electrons to break free from the atom, and the rest is left over as kinetic energy of the ejected electron. • Threshold energy is required *

31. * AP* • A nitrogen gas laser pulse with a wavelength of 337 nm contains 3.87 mJ of energy. How many photons does it contain?

32. Problems So far… • Rutherford suggested that electrons move around the nucleus as planets orbit the Sun. • Recall: accelerating objects change speed and/or direction. Electron travelling in circular orbit is constantly changing it’s directions- therefore acceleration • Orbiting electrons should emit photos of electromagnetic radiation, losing energy, spiraling towards the nucleus and collapsing the atom. NOT TRUE – atoms are stable.

33. Bright Line & Dark Line Spectra • Robert Bunsen & GustvanKirchoff invented the spectroscope ( method for analyzing spectra) • Discovered a bright-line spectrum, a series of bright lines produced or emitted by an excited gas ( documented, new spectra= new element) • Cesium and Rubidium were discovered like this!

34. Remember Grade 10? (Neither do I !) • When white light is shone through a prism, it is broken into a spectrum. • Each colour corresponds to a different wavelength of light. • Each wavelength corresponds to a particular amount of energy.

36. Emission- Spectra The pattern of lines emitted by excited atoms of an element is unique = atomic emissionspectrum

37. Absorption or dark-line spectra were discovered in 1814. They were a series of dark lines (missing parts), of a continuous spectrum, produced by placing a gas between the continuous spectrum source and the observed.

38. Bohr’s Model of the Atom • Neils Bohr suggested that understanding Thomson’s work with cathode rays required the use of the new quantum theory of light proposed by Planck and Einstein. • Bohr and Thomson argued all the time … so Bohr abandoned him … and went to work with Rutherford instead (in Manchester!) Rutherford also abandoned Thomson … hmmm…..

39. Bohr’ Model • Bright & Dark-light emission spectra meant that only certain quanta (photon energies) can be released or absorbed • Therefore, the energy of the electron inside the atom must also be quantized. • Therefore, electrons can only have certain energies

40. Bohr’s First Postulate • Electrons do not radiate energy as they orbit the nucleus. Each orbit corresponds to a state of constant energy ( called a stationary state)

41. Bohr’s Second Postulate • Electrons can change their energy only by undergoing a transition from one stationary state to another. • High-to-low-energy-state = emitting energy (bright line spectrum) • Low-to-high-energy-state = absorbing energy (dark-line spectrum)

42. Successes and Weaknesses of Bohr’s Model • Successes • Bohr’s mathematics explained all the observations for Hydrogen perfectly • This is a major success for Quantum Mechanics – to this day it has never failed • He solved Rutherford’s problem • Weaknesses • Bohr’s method only worked for Hydrogen • Although the quantum theory of light  was experimentally proven, other experiments had proven that light had also continuous  wavelike properties. Einstein suggested that there were "two contradictory pictures of reality; separately neither of them fully explains the phenomena of light, but together they do". Hence light and photons display wave and particle properties

43. *AP* En = –B/n2 where B is a constant = 2.179 × 10–18 J and n is an integer

44. *AP* • En = –B/n2 where B is a constant = 2.179 × 10–18 J and n is an integer • B = hcRh (h = planck’s constant, speed of light, Rydberg constant which is 1.096776 x 107 m-1. • The lower ( more negative) the energy is, the more stable the atom will be. This happens for n=1. • As n gets larger, the energy becomes less negative, and therefore increases.

45. For the final and initial levels: The energy difference between nf and ni is: *AP* Bohr’s equation is most useful in determining the energy change (Elevel) that accompanies the leap of an electron from one energy level to another

46. *AP* If an electron moves from n1 = 3, to nf=1 …..