Atomic Structure

1. Atomic Structure Unit 2

2. Overview • Atomic Theory • John Dalton • Law of Conservation of Mass • Law of Definite Proportions • Law of Multiple Proportions • Ernest Rutherford • Robert Millikan • J.J. Thompson • Atomic Structure • Protons, neutrons, electrons • Atomic number • Isotopes • Mass number • Average atomic mass • Wave nature of light • Electromagnetic Spectrum • C = λv • Bohr Models • Photoelectric effect • Absorption/emission • E = hc/ λ • Heisenberg Uncertainty Principle • Quantum numbers • Pauli Exclusion Principle • Hund’s Rule • Aufbau Principle • Configurations (orbital, electron, noble gas) • Paramagnetism/diamagnetism • Exceptions

3. Chemistry TimeLine B.C. 400 B.C. Democritus and Leucipposuse the term "atomos” 2000 years of Alchemy • 1500's • Georg Bauer: systematic metallurgy • Paracelsus: medicinal application of minerals 1600's Robert Boyle:The Skeptical Chemist.Quantitative experimentation, identification of elements • 1700s' • Georg Stahl: Phlogiston Theory • Joseph Priestly: Discovery of oxygen • Antoine Lavoisier: The role of oxygen in combustion, law of conservation of mass, first modern chemistry textbook

4. Chemistry timeline • 1800's • Joseph Proust: The law of definite proportion (composition) • John Dalton: The Atomic Theory, The law of multiple proportions • Joseph Gay-Lussac: Combining volumes of gases, existence of diatomic molecules • Amadeo Avogadro: Molar volumes of gases • JonsJakob Berzelius: Relative atomic masses,modern symbols for the elements • Dmitri Mendeleyev: The periodic table • J.J. Thomson: discovery of the electron • Henri Becquerel: Discovery of radioactivity • 1900's • Robert Millikan: Charge and mass of the electron • Ernest Rutherford: Existence of the nucleus, and its relative size • Meitner & Fermi: Sustained nuclear fission • Ernest Lawrence: The cyclotron and trans-uranium elements

5. The Greeks • 400 BC • Democritus • Matter consists of small particles • Called them “atomos” • Idea rejected by peers • No scientific proof

6. The Greeks (cont…) • Aristotle • All matter continuous • 4 elements = earth, water, air, and fire • No scientific proof • Idea endured for 2000 years

7. John Dalton - 1808 • School Teacher • Atomic Theory • All matter is composed of extremely small particles called atoms. There are different kinds called elements. • Atoms of the same element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. • Atoms cannot be subdivided, created, or destroyed. • Atoms of different elements combine in simple, whole number ratios to form chemical compounds. • In chemical reactions, atoms are combined, separated, or rearranged but never destroyed/created.

8. Laws derived from dalton • Law of Conservation of Mass • Total mass present before chemical reaction is same as mass after chemical reaction • 2H2O  2H2 + O2 If you have 10 grams of water to start, you will get 1.12 g of hydrogen and 8.88 g of oxygen • Law of Constant Composition (definite proportions) • Relative numbers and kinds of atoms are constant • Water is 88.8% oxygen and 11.2% hydrogen by mass no matter how much you have • Law of Multiple Proportions • If two elements combine to form more than one compound, the masses of the two elements are in the ratio of small whole numbers • CO2 versus CO (mass ratio is 2 to 1 for oxygen)

9. J.J. Thomson • British Physicist • Discovered electron • Cathode-ray experiment • Plum pudding view of atom

10. Thompson Cathode Ray Experiment • Electric current sent through gases in glass tube called cathode-ray tube • Surface of tube opposite the cathode glowed – caused by stream of particles • Ray traveled from cathode to anode • Cathode rays deflected by magnetic field away from negatively charged object (like a magnet) • Cathode rays concluded to have negative charge

11. Robert Millikan - 1909 • American Physicist • Charge on each electron is same • Charge of electron is -1.6022 x 10-19C • Calculated mass of electron as 9.10x 10-31 kg • Oil drop experiment

12. Millikan Oil Drop Experiment • Drops of oil that had picked up extra electrons allowed to fall between two electrically charged plates • Measured how voltage on plates affected rate of fall • Calculated charges of drops then deduced charge of a single electron on the drops

13. Ernest Rutherford • Discovered nucleus • Planetary model of the atom

14. Rutherford Gold Foil Experiment • Bombarded thin piece gold foil with alpha particles (positively charged particle 4 times mass of hydrogen atom) • Expected to pass right through gold foil • 1 in 8000 particles deflected back toward source • “As if you fired 15-inch artillery shell at a piece of tissue paper and it came back and hit you” • Concluded most of atom is empty space except for a very small force within atom • Called positive bundle of matter the “nucleus”

15. Modern Atomic Theory • Atom consists of proton, neutron, and electron • Proton charge = +1 • Neutron charge = 0 (neutral) • Electron charge = -1 • Protons and Neutrons located in nucleus • 99.9% of atom’s mass is in nucleus • Electrons located outside the nucleus

16. 47 Silver Ag 107.87 Element Blocks Atomic number Name of the element Element Symbol Atomic mass

17. Element Blocks • Atomic Number • equal to number of protons in an atom • Element Symbol • First letter always capitalized • If second letter exists, it is lowercase

18. isotopes • Isotopes are atoms of the same element having different masses due to varying numbers of neutrons.

19. Atomic Mass • Atomic mass is the average of all the naturally isotopes of that element. Carbon = 12.011

20. Mass Number • Mass Number = Protons + Neutrons • Not found on periodic table • Isotopes have different mass numbers (due to neutrons)

21. Symbolizing Elements C– 12 Atomic number Mass number Mass number

22. Wave-Particle Duality • JJ Thomson won the Nobel prize for describing the electron as a particle • His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron. The electron is a particle! The electron is an energy wave!

23. Traveling Waves Much of what has been learned about atomic structure has come from observing the interaction of visible light and matter.

24. Wave Theory of Electron • 1924De Broglie suggested that electrons have wave properties to account for why their energy was quantized. • He reasoned that the electron in the hydrogen atom was fixed in the space around the nucleus. • He felt that the electron would best berepresented as a standing wave. • As a standing wave, each electron’s path must equal a whole number times the wavelength.

25. De Broglie The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves. Louis deBroglie

26. Waves • Wavelength, l • The distance for a wave to go through a complete cycle. • Amplitude • Half of the vertical distance from the top to the bottom of a wave. • Frequency, n • The number of cycles that pass a point each second.

27. Waves

28. Waves • Longer wavelength = lower frequency = lower energy • Shorter wavelength = higher frequency = higher energy

29. Wavelength Frequency Relationship • The SI unit of frequency (n) is the hertz, Hz 1 Hz = 1 s-1 • Wavelength and frequency are related c = ln c is the speed of light, 2.998 x108 m/s

30. Practice Problem The wavelength of an argon laser's output is 488.0 nm. Calculate the frequency of this wavelength of electromagnetic radiation. c = ln • Convert nm to m 488 nm x (1 m / 109 nm) = 4.88 x 10-7 m • Then, substitute into c = λν (4.88 x 10-7 m) (v) = 3.00 x 108 m s-1 v = 6.15 x 1014 s-1 = 6.15 x 1014 Hz

31. Electromagnetic Radiation • Electromagnetic Radiation • Energy in the form of transverse magnetic and electric waves. • Electromagnetic Spectrum • Contains all forms of electromagnetic radiation • Visible spectrum • Portion of electromagnetic spectrum that we can see (colors)

32. Electromagnetic Spectrum

33. Separation of Light • ‘White’ light is actually a blend of all visible wavelengths. They can separated using a prism.

34. Line Spectra • Neils Bohr studied the spectra produced when atoms were excited in a gas discharge tube.

35. Line Spectra • Each element produces its own set of characteristic lines

36. Bohr Model • Bohr proposed a model of how electrons moved around the nucleus. • He wanted to explain why electrons did not fall in to the nucleus. • He also wanted to account for spectral lines being observed. • He proposed that the energy of the electron was quantized - only occurred as specific energy levels.

37. BoHr Model • In the Bohr model, electrons can only exist at specific energy levels (orbit). • Each energy level was assigned a principal quantum number, n. Energy

38. Bohr Model • The Bohr model is a ‘planetary’ type model. • Each principal quantum represents a new ‘orbit’ or layer. • The nucleus is at the center of the model.

39. Transitions Electron transitionsinvolve jumps of definite amounts Of energy.

40. Absorption Emission • Absorption – Electromagnetic radiation is absorbed by an atom causing electrons to jump to a higher energy state (excited state). • Emission – Energy is released by an atom as particle of light (photon) as electrons fall back to the lower energy state (ground state). • Depending on frequency of photon, different colored light may be seen

41. Particle Properties • Although electromagnetic radiation has definite wave properties, it also exhibits particle properties. • Photoelectric effect. • First observed by Hertz and then later explained by Einstein. • When light falls on Group IA metals, electrons are emitted (photoelectrons). • As the light gets brighter, more electrons are emitted. • The energy of the emitted electrons depends on the frequency of the light.

42. Photoelectric Effect • The energy of a photon is proportional to the frequency. (Photon energy) E= hn • The energy is inversely proportional to the wavelength (remember c =λν so v =c/λ ). E = hc /l h is Plank’s constant, 6.626 x 10-34 J .S c is the speed of light, 2.998 x108 m/s

43. Photon Energy Example • Determine the energy, in kJ/mol of a photon of blue-green light with a wavelength of 486 nm. • E = • = • = 4.09 x 10-19 J h c l (6.626 x 10-34 J.s)(2.998 x 108 m.s-1) (4.86 x 10-7 m)

44. h mv l = De Broglie Equation • l = wavelength, meters • h = Plank’s constant • m = mass, kg • v = frequency, m/s

45. h mv 6.6 x 10-34 kg m2 s-1 (9.1 x 10-31 kg)(2.2 x 106 m s-1) l = De Broglie Equation • Using De Broglie’s equation, we can calculate the wavelength of an electron. l = = 3.3 x 10-10 m The speed of an electron had already been reported by Bohr as 2.2 x 106 m s-1.

46. Heisenberg Uncertainty PRinciple • In order to observe an electron, one would need to hit it with photons having a very short wavelength. • Short wavelength photons would have a high frequency and a great deal of energy. • If one were to hit an electron, it would cause the motion and the speed of the electron to change. • According to Heisenberg, it is impossible to know both the position and the speed of an object precisely.

47. Quantum Model • Schrödinger developed an equation to describe the behavior and energies of electrons in atoms. • His equation is similar to one used to describe electromagnetic waves. • Each electron can be described in terms of its quantum numbers.

48. Quantum Numbers • Each electron in an atom has a unique set of 4 quantum numbers which describe it. • Principal quantum number • Angular momentum quantum number • Magnetic quantum number • Spin quantum number

49. Quantum Numbers • Principal quantum number, n • Tells the size of an orbital and largely determines its energy. n = 1, 2, 3, ……

50. Quantum Numbers • Angular momentum, l • The number of subshells that a principal level contains. It tells the shape of the orbitals. l= n – 1 to 0 • Orbitals • An orbital is a region within an energy level where there is a probability of finding an electron • Orbital shapes are defined as the surface that contains 90% of the total electron probability.