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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 Configurations (orbital, electron, noble gas) Pauli Exclusion Principle Hund’s Rule 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. Atomic mass unit (AMU) Too difficult to measure elements in “grams” so we use the atomic mass unit Approximately the mass of 1 proton or 1 neutron Relative to the carbon atom 1 amu is 1/12 the mass of the carbon atom
19. isotopes Isotopes are atoms of the same element having different masses due to varying numbers of neutrons.
20. Atomic Mass Atomic mass is the average of all the naturally isotopes of that element. Carbon = 12.011
21. Mass Number Mass Number = Protons + Neutrons Not found on periodic table Isotopes have different mass numbers (due to neutrons)
22. Symbolizing Elements C– 12 Mass number Mass number Atomic number
23. 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!
24. Traveling Waves Much of what has been learned about atomic structure has come from observing the interaction of visible light and matter.
25. 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.
26. 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
27. 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.
28. Waves
29. Waves Longer wavelength = lower frequency = lower energy Shorter wavelength = higher frequency = higher energy
30. 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
31. 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
32. 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)
33. Electromagnetic Spectrum
34. Separation of Light ‘White’ light is actually a blend of all visible wavelengths. They can separated using a prism.
35. Line Spectra Neils Bohr studied the spectra produced when atoms were excited in a gas discharge tube.
36. Line Spectra Each element produces its own set of characteristic lines
37. 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.
38. 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
39. 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.
40. Transitions Electron transitionsinvolve jumps of definite amounts Of energy.
41. 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
42. 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.
43. 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
44. 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)
45. h mv l = De Broglie Equation l = wavelength, meters h = Plank’s constant m = mass, kg v = frequency, m/s
46. 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.
47. 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.
48. 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.
49. Quantum Numbers Each electron in an atom has a unique set of 4 “numbers” which describe it Energy level Orbital shape Orientation Spin
50. Main Energy level Principal quantum number, n Tells the size of an orbital and largely determines its energy. n = 1, 2, 3, ……
51. Shape of the orbital Angular momentum The number of subshells that a principal level contains. It tells the shape of the orbitals. s p d f 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.
52. Orbital Shapes
53. Orientation of orbital Magnetic quantum number, ml Describes the direction that the orbital projects in space. Think in terms of axes “x, y, z”
54. Spin of the electron Pauli added one additional quantum number that would allow only two electrons to be in an orbital. Spin quantum number, ms. An electron can spin clockwise or counterclockwise
55. Rules Pauli exclusion principle Pauli proposed that no two electrons in an atom can have the same set of four quantum numbers Unless you live together, none of you have the exact same address
56. Other Rules Aufbau Principle Electrons are placed into orbitals, subshells, and shells in order of increasing energy
57. Other Rules Hund’s Rule The most stable arrangement of electrons in a subshell is the one in which electrons have the most number of parallel spins possible.
58. Orbital notation of electrons Graphical representation of an electron configuration One arrow represents one electron Shows spin and which orbital within a sublevel Follow all rules(Aufbau principle, two electrons in each orbital, etc. etc.)
59. Orbital notation Use atomic number as number of electrons in an atom He Be Mg Si Ne
60. Magnetism Diamagnetism Elements have all of their electrons spin paired All of an element’s subshells are completed Not affected by magnetic fields Paramagnetism Not all electrons are spin paired in an element Most elements are this Affected by magnetic fields
61. Electron Configuration A list of all the electrons in an atom (or ion) Must go in order (Aufbau principle) 2 electrons per orbital, maximum We need electron configurations so that we can determine the number of electrons in the outermost energy level. These are called valence electrons. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14…etc.
62. Electron configuration 2p4 Number of electrons in the sublevel Energy Level Sublevel
63. Examples He, 2: 1s2 Ne, 10: 1s2 2s2 2p6 Ar, 18: 1s2 2s2 2p6 3s2 3p6 Kr, 36: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
64. Periodic Table Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental) s orbitals d orbitals p orbitals f orbitals
65. Why are d and f orbitals always in lower energy levels d and f orbitals require LARGE amounts of energy It’s better (lower in energy) to skip a sublevel that requires a large amount of energy (d and f orbtials) for one in a higher level but lower energy
66. Noble Gas Notation A way of abbreviating long electron configurations Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ]. Step 2: Find where to resume by finding the next energy level. Step 3: Resume the configuration until it’s finished. Example: [Ne] 3s2 3p5
67. Exceptions Remember d and f orbitals require LARGE amounts of energy If we can’t fill these sublevels, then the next best thing is to be HALF full (one electron in each orbital in the sublevel) There are many exceptions, but the most common ones are For the purposes of this class, we are going to assume that ALL atoms (or ions) that end in d4 or d9 are exceptions to the rule. This may or may not be true, it just depends on the atom.
68. Exceptions d4 is one electron short of being HALF full In order to become more stable (require less energy), one of the closest s electrons will actually go into the d, making it d5 instead of d4. For example: Cr = [Ar] 4s2 3d4 Since this ends exactly with a d4 it is an exception to the rule. Thus, Cr = [Ar] 4s1 3d5 Remember, half full is good… and when an s loses 1, it too becomes half full! d9 works the same way