Chemistry I Honors Unit 3: Quantum Mechanical Model of the Atom & Periodic Trends

1. Chemistry I HonorsUnit 3: Quantum Mechanical Model of the Atom & Periodic Trends

2. Objectives #1-7: The Development of a New Atomic Model • I. Electromagnetic Radiation • In the late 1800’s and early 1900’s, scientists discovered that passing an electric current through gases of various elements caused electromagnetic radiationin the form of colored lightto be emittedfrom the gas.

3. Examples of colors produced by the electrically charged gasses include:

4. Additional testing showed that EMR of energies too low or too high to see with the eye were also produced • Electromagnetic radiation is energy that travels in the form of a wave

5. Examples of Electromagnetic Waves

6. Characteristics of EMR: • The wavelength of a wave is the distance between the peaksof the wave • The frequency of a wave is the number of peaks that pass by a point is space in one second (the rate of reproducibility) • The speed of all EMR is the speed of light(3.00 X 108 m/s)

7. Wavelength and frequency are inverselyrelated to each other c = (f) () • Frequency and energy are directly related to each other E = (h) () • Wavelength and energy are inversely related to each other E = (h)(c) / ()

8. Types of EMR Lower Energy Higher Energy Radio Radar Micro IR Visible* UV X-rays Gamma red, orange, yellow, green, blue, indigo, violet

9. Objectives #1-7: The Development of a New Atomic Model • Electromagnetic Radiation * EMR refers to all the various types of radiant energy, from radio waves to gamma waves

10. II. The Origins of Wave MechanicsEMR has dualqualities: (Louis DeBroglie, 1892-1987), French EMR is like 4 year old—it can’t make up its mind what it wants to be for Halloween!! “A wave! No, wait, a particle! No wait, a wave… no wait…..”

11. EMR acts as a particle when it interacts with matter; this is illustrated by the photoelectric effect which involves the emission of electrons when radiation of a specific frequency strikes the surface of a metal(Albert Einstein, 1905, German-American)  = (h) (v0) • It acts as a wave when it travels through space •  = (h) /(m) (v)

12. Albert Einstein, 1905, German-American The photoelectric effect helps explain why your solar calculator works!

13. Photoelectric Effect

14. Wave Particle Duality  = (h) / (m) (v)

15. Objectives #1-7 The Development of a New Atomic Model • Atomic Emission Spectrum: an explanation of the colors produced by exciting atoms • Niels Bohr, 1885-1962, Danish

16. Examples of Atomic Emission Spectrums

17. Bohr’s Theory of Light Emission • An electron is normally in its low energy state or ground state. • When the electron becomes excitedwith a certain amount of energy or quantum, it will “jump” to a higher level of energy or its excited state. • This new state is unstable for the electron and so this excess energy is emitted as a photonof EMR and the electron returns to the ground state.

18. The Bohr Model and Electron Transitions Illustrated

19. The Bohr Model of the Atom • Proposed that electrons revolve around the nucleus in definite paths or orbits • Each electron has a certain amount of energy associated with it • Electrons are confined to specific energy levels • In order to move from one level to the next, an electron must absorb or release a certain quantumof energy

20. Operation of Spectroscope

21. The Bohr Model Let’s draw some Bohr diagrams!!!!

22. IV. The Quantum Mechanical Model of the AtomErwin Schrodinger, 1887-1961, Austrian • “Matter also has particle and wave characteristics. So since matter is made of atoms, and atoms contain electrons, electrons could also travel in waves.”

23. Proposed the quantum mechanical model for electrons. • The exactpath of the electron can not be determined because it is traveling near the speed of light and is too small in size. • This idea was based on the work of Werner Heisenberg, 1901-1976, German. • In the Heisenberg Uncertainty Principle there is a limit to how certain we can be about the position and speed of very tiny particles such as electrons.

24. Werner Heisenberg, 1901-1976, German

25. Evolution of the Bohr Model into the Quantum Mechanical Model

26. Heisenberg Uncertainty Principle

27. In the quantum model only the probability of finding the electron in a certain area can be determined. • The most highly probable location for an electron about the nucleus is the orbital. • The combination of these areas about the nucleus is called the electron cloud.

28. Objectives #8-10: Quantum Numbers • Schrodinger’s Equation: EΨ = -h2/2m(ð2Ψ/ðx2 + ðΨ/ðy2 + ð2Ψ/ðz2) + V(x, y, z)Ψ (Neat, huh? No you do NOT have to use it, memorize it or solve it!!! ) • Solving the previous equation produces various orbital shapes, just as solving y = 1/2x + 2 produces a straight line.

29. Quantum Numbers • Describes energy and location of electrons • Every electron in an atom is unique; each electron has a different energy and therefore will have a different set of quantum numbers.

30. Principle Quantum Number (n) • Indicates energyand distance from nucleus • Indicates energy level number • Can take on values of: 1 → infinity • Orbital (Angular Momentum) Quantum Number (l) • Indicates shape of orbital (sublevel) • Can take on values of: n-1 (0, 1,2,3, etc) • Possible orbital shapes:

32. Orbital Shapes

33. The number of orbital (sublevel) shapes in a level is equal to the level number:

34. C. Magnetic Quantum Number (ml) • Indicates the orientation (direction) of the orbital in space • Indicates the number of orbital directions in a sublevel • Can take on values of: 0 and –l  +l

35. C. Magnetic Quantum Numbers (ml)

36. Spin Quantum Number (ms) • Indicates the direction of electron spin • Can take on values of: +1/2, -1/2 • No more than 2 electrons can occupy a single orbital • Problems Involving Quantum Numbers • See notes

37. IV. Summary of Electron Energy Level Capacities (See Chart in Lecture Guide) • Some relationships to notice: • If “n” is the number of levels, then the number of sublevels is equal to “n” • If “n” is the number of levels, then the total number of orbitals in a level is equal to n2 • If “n” is the number of levels (and every orbital can hold up to 2 electrons), then the total number of electrons in a level is equal to 2n2

38. Objectives #11-12: Electron Configurations • Electron configurations show electron arrangement • Rules Governing Electron Configurations • The Aufbau Principle • Electrons enter orbitalsof lowest energy first • The Pauli Exclusionary Principle • An atomic orbital may describe, at most, 2electrons

39. Hund’s Rule • Electrons enter orbitals of the same energy with the samespin until each orbital contains one electron before pairing begins • Examples of Electron Configurations w/Orbital Notations & Noble-gas Configurations • H, He, Li, C, P • Diagonal Rule • More Examples: • Noble-gas configurations: • Exceptions to the Aufbau Principle

40. Objectives #13-20 The Periodic Table / Periodicity of Properties • Development of the Periodic Table • The work of Mendeleev (1871, Russian) • Elements were grouped by their properties; • Allowed for prediction of new elements; • Elements arranged by increasing atomic mass • The work of Moseley (1911, English): • Used X-ray studies to determine atomic numbers of elements; • Arranged in order of increasing atomic number

41. Organization of Modern Periodic Table • See PT handout III. Valence and the Periodic Table • In most chemical reactions only the valenceelectrons are involved • Valence electrons are the outer level electrons in an atom • Examples—see lecture guide

42. Objectives #13-20 The Periodic Table / Periodicity of Properties • Periodic Properties • Periodic Law: The properties of the elements vary periodically—in a predictable pattern— when placed in order of increasing atomic numbers • Periodic Trends • Atomic Radii • 1/2 the distance between the nuclei of identical atoms that are bonded together (how big the atom is!)

43. Objectives #13-20 The Periodic Table / Periodicity of Properties • Increases down a group • Decreases across the period • A cation or positive ion is smallerthan its original parent atom • An anion or negative ion is larger than its original parent atom • Ionization Energy • The amount of energy required to remove an electron from an atom

44. Objectives #13-20 The Periodic Table / Periodicity of Properties • decreases down a group • increases across the period • Electronegativity • A measure of the ability of an atom to gainelectrons during the bondingprocess • decreases down a group • increasesacross the period

45. Objectives #13-20 The Periodic Table / Periodicity of Properties • Why these variations occur: • Adding additional energy levels increases size of atom • The shielding effect (inner energy levels block influence of nucleus from outer energy levels) also increases • Increasing nuclear charge holds electrons slightly closer to the nucleus across period