Unit 7: Electronic Structure and the Periodic Table

1. Some people succeed because they are destined to, but most people succeed because they are determined to.

2. Unit 7:Electronic Structure and the Periodic Table Chemistry The GowerHour

3. Review Democritus John Dalton A. Greek philosopher, ___________, originated the idea of the atom. (atomos) B. About 1800, ___________ developed the Atomic Theory. C. J.J. Thompson discovered the ________, and wrote the _________ model of the atom. • __________ discovered the proton and the nucleus using alpha particles and the gold foil experiment. He created the ______________ of the atom. E. Neils _____ created a model of the atom that explained the atomic line spectrum of _________. F. The _______________________ (our current model) explained the line spectra of all the atoms. This model relates light to electron movement within the orbitals. electron raisin bun Rutherford planetary model Bohr hydrogen quantum mechanical model

4. elements I. The Development of Atomic Models A. Rutherford’s atomic model could not explain the chemical properties of _________. B. Rutherford’s atomic model could not explain why objects change ______ when heated (Atomic line spectra). color

5. Page 128 C. Timeline 1. 1803 to 1911.

6. 2. 1913 to 1932

7. specific D. The Bohr Model 1. Bohr proposed that an electron is found only in ________ circular paths, or orbits, around the nucleus. 2. Each possible electron orbit in Bohr’s model has a fixed energy. • The fixed energies an electron can have are called _____________. • A _________ of energy is the amount of energy required to move an electron from one energy level to another energy level. energy levels quantum

8. 5.1 3. Like the rungs of the strange ladder, the energy levels in an atom are ____ equally spaced. 4. The higher the energy level occupied by an electron, the ____ energy it takes to move from that energy level to the next higher energy level. not less

9. likely E. The Quantum Mechanical Model 1. The quantum mechanical model determines the allowed energies an electron can have and how _______ it is to find the electron in various locations around the nucleus. 2. Austrian physicist Erwin Schrödinger (1887–1961) used new theoretical calculations and results to devise and solve a mathematical equation describing the behavior of the electron in a hydrogen atom. 3. The modern description of the electrons in atoms, the __________________________, comes from the mathematical solutions to the Schrödinger equation. quantum mechanical model

10. probability 4. The propeller blade has the same ___________ of being anywhere in the blurry region, but you cannot tell its location at any instant. The electron cloud of an atom can be compared to a spinning airplane propeller.

11. more dense 5. In the quantum mechanical model, the probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloud. The cloud is ____________ where the probability of finding the electron is high. 90 % probability

12. 5.1 atomic 6. An ________ orbital is often thought of as a region of space in which there is a high probability of finding an electron. 7. Each energy sublevel corresponds to an orbital of a different _______, which describes where the electron is likely to be found. shape

13. 5.1 spherical dumbbell 8. Different atomic orbitals are denoted by letters. The s orbitals are __________, and p orbitals are __________-shaped. 9. Four of the five d orbitals have the same shape but different orientations in space.

14. probable n l F. Quantum Numbers: Describe the most ________________ location of the electron. 1. Principle Energy Level (___) n = 1, 2, 3, 4, … 2. Sublevel (____): Shape of orbital Shapes: s, p, d, f y y x x z z l = 0 l = 1 s p y y x x z z l = 3 l = 2 f d

15. 0 -2 -1 0 +1 +2 -1 0 +1 p s d -3 -2 -1 +1 +2 +3 0 f l = 0 (s orbitals) ml 3. Orbital (orientation in space) (____) -1 0 + 1

16. ms arrow 4. Spin of electron (____) Two electrons in the same orbital must have opposite spins. This produces a magnetic field in opposite directions. We represent a magnetic field by using an _________. G. Address Analogy n = energy level l = sublevel ml = orbital ms = spin ________ _________ _______ _______ ms = +½ ms = -½ state city street house

17. H. Number of Orbitals per Subshell 2 s 1 3 p 6 d 10 5 f 14 7

18. I. Allowed sublevels in each energy level: n Allowed sublevels 1 1s 2 2s, 2p 3 3s, 3p, 3d 4 4s, 4p, 4d, 4f 5 5s, 5p, 5d, 5f, 5g

19. electrons lowest II. Electron Configurations: The arrangement of ________ around the nucleus.A. Rules for electron configuration A. Aufbau Principle: Electrons fill the _________ orbitals first. (Lowest energyto highest energy).  B. Pauli Exclusion Principle: No two electrons can have four __________ quantum numbers, therefore an orbital can hold no more than ____ electrons of opposite _____(represented by ____ and _______ arrows). C. Hund’s Rule: When electrons fill __________ of equal energy, they fill the orbitals one at a time until they have to pair up. Each seat is uniquely identified (E, R12, S8) Each seat can hold only one individual at a time identical spin up two down orbitals

20. number of electrons in the orbital principal quantum number n Sublevel Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom. 1s1 Orbital diagram H 1s

21. ? ? “Fill up” electrons in lowest energy orbitals (Aufbau principle) Li 3 electrons Be 4 electrons B 5 electrons C 6 electrons B 1s22s22p1 Be 1s22s2 Li 1s22s1 H 1 electron He 2 electrons He 1s2 H 1s1

22. s block 1s 1 2 D. Using the Periodic Table to determine electron configurations: ***Label your periodic table with the energy subshells***

23. Fig. 7.27

24. 2p5 1s2 2s2 1s 2s 2p 2p6 1s2 2s2 3s2 3p2 1s 2s 2p 3p 3s 2s2 2p6 3s2 3p6 1s2 4s2 3d4 3s 3s 1s 1s 2s 2s 2p 2p 3p 3p 4s 4s 3d 3d 2p6 3s2 3p6 1s2 4s2 3d10 4p4 2s2 4p 9 E. Examples using the Periodic Table Element# e-e- ConfigurationOrbital Diagram 1. F 2. Si 3. Cr 4. Se 14 24 34

25. [Ne] 3s2 3p2 [Ne] 3s 3p [Ar] 4s2 3d4 [Ar] 4s 3d [Ar] 4s2 3d10 4p4 [Ar] 4s 3d 4p 4d1 [Kr] 5s2 [Kr] 5s 4d [Xe] 6s2 4f 14 5d8 [Xe] 6s 4f 5d Noble gas 14 F. Kernal Configurations (Abbreviated Electron Configurations): Use the preceding ________. 1. Si 2. Cr 3. Se 4. Y 5. Pt 24 34 39 78

26. 1s2 2s2 2p6 3s1 1s2 2s2 2p6 1s2 2s2 2p6 3s2 3p1 1s2 2s2 2p6 Add 1 e- Lose 1 e- 1s2 2s2 2p5 1s2 2s2 2p6 Add 3 e- Lose 3 e- 1s2 2s2 2p3 2s2 2p6 1s2 Ne Ne Ne Ne 1s2 1s2 1s2 1s2 1s2 2s2 2s2 2s2 2s2 2s2 2p6 2p6 2p6 2p6 2p6 G. Electron Configurations of Ions: 1. Cations: Na Na+ Al Al3+ 2. Anions: F F- N N3- Ne 3. Isoelectronic: Elements and ions with the _____ electron configurations. isoelectronic same

27. [Ar] [Ar] 4s2 4s2 3d9 3d4 [Ar] 4s1 3d5 [Ar] 4s1 3d10 [Kr] 5s1 4d10 [Xe] 6s1 4f 14 5d10 d H. Irregularities with the Transition Metals: Half-filled and filled __ subshells are more stable than other configurations, so some atoms have irregular electron structures. 1. Examples: Chromium Copper Silver Gold (half-filled) (full d)

28. s p n I. Valence Electrons: The electrons in the outermost __ and __ subshells. 1. The value of _____ must be the same for the s and p. 2. The number is never greater than ___ (____) 3. Elements can be represented with the dot formula: Li Be B C N O F Ne 8 s2p6

29. [Ar] [Ar] 3d5 3d6 [Ar] 4s2 3d6 [Kr] 5s2 4d10 5p2 [Kr] 5s2 4d10 [Kr] 4d10 J. Transition Metals Lose e- from valence electrons first!! e- configuration of transition metal cations Ex. Fe Fe2+ Fe3+ Sn Sn2+ Sn4+

30. s subshell p subshell 1st Energy Level 2nd Energy Level 1s 2s 2p p orbitals n Subshell Review of Quantum Model 10 Ne n2 = # orbitals in energy level 2n2 = # e- in energy level 1 1s 2 2s 2p 3 3s 3p 3d 4 4s 4p 4d 4f 5 5s 5p 5d 5f 5g

31. 5.3 amplitude III. Light is a wave A. The __________ of a wave is the wave’s height from zero to the crest. B. The __________, represented by  (the Greek letter lambda), is the distance between the crests. C. The __________, represented by  (the Greek letter nu), is the number of wave cycles to pass a given point per unit of time. D. The SI unit of cycles per second is called a hertz (Hz). wavelength frequency

32. 5.3 E. The wavelength and frequency of light are inversely proportional to each other.

33. 5.3 F. The product of the frequency and wavelength always equals a constant (c), the speed of light. The wave equation: G. According to the wave model, light consists of ______________ waves. H. Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays. I. All electromagnetic waves travel in a vacuum at a speed of 2.998  108 m/s. (3.00 x 108 m/s) electromagnetic

34. 5.3 J. Sunlight consists of light with a continuous range of wavelengths and frequencies. K. When sunlight passes through a prism, the different frequencies separate into a ________ of colors. L. In the visible spectrum, ____ light has the longest wavelength and the lowest frequency. _______ light has the shortest wavelength and the highest frequency. M. The electromagnetic spectrum consists of radiation over a broad band of wavelengths. spectrum red Violet

35. 5.3

36. 1. Example: The green light associated with the aurora borealis is emitted by excited (high-energy) oxygen atoms at 557.7 nm. What is the frequency of this light? = 557.7 x 10 -9 m c = 3.00 x 10 8 m/s f = ?

37. f = 4.14 x 10 14 s-1 c = 3.00 x 10 8 m/s 2. Example: If red light has a frequency of 4.14 x 1014 Hz, what is its wavelength? 3. Example: What is the wavelength (in meters) of an FM radio station that has a frequency of 101.5 MHz? = 725 nm f = 101.5 MHz = 101.5 x 10 6 Hz c = 3.00 x 10 8 m/s

38. IV. Light is a particle photons 1. Light is a stream of particles called _______, with energy given by the equation: E = h = • Energy is ________ related to wavelength.  ___ E ___ • Energy is ______ related to frequency.  ___ E ___ Energy (J) Planck’s constant = 6.626 x 10-34 J·s inversely ↑ ↓ directly ↑ ↑

39. c = 3.00 x 10 8 m/s = 6.1 x 1014 Hz h = 6.626 x 10 -34 Js 4. Calculate (a) the frequency, and (b) energy of a monochromatic blue light that has a wavelength of 490 nm. 5. If a photon has an energy of 7.55 x 10-19 J, what is its wavelength in nanometers? h = 6.626 x 10 -34 Js c = 3.00 x 10 8 m/s = 263 nm

40. IQ1: Review Problems f = 1090 KHz = 1090 x 10 3 Hz c = 3.00 x 10 8 m/s h = 6.626 x 10 -34 Js 1. Example: Calculate (a) the wavelength, and (b) the energy of a radiowave with a frequency of 1090 kHz (AM radio). E = ? = 275 m E = 7.22 x 10 -28 J

41. higher emitting V. Atomic Spectra A. When atoms absorb energy, electrons move into ______ energy levels. These electrons then lose energy by ________ light when they return to lower energy levels. B. A prism separates light into the colors it contains. When white light passes through a prism, it produces a ________ of colors. rainbow

42. 5.3 discrete C. When light from a helium lamp passes through a prism, ________ lines are produced.

43. 5.3 spectrum D. The frequencies of light emitted by an element separate into discrete lines to give the atomic line _________ of the element. Mercury Nitrogen

44. E. An Explanation of Atomic Spectra 1. In the Bohr model, the lone electron in the hydrogen atom can have only certain specific energies. a. When the electron has its lowest possible energy, the atom is in its _____________. b. Excitation of the electron by absorbing energy raises the atom from the ground state to an _______ state. c. A ________ of energy in the form of light is emitted when the electron drops back to a lower energy level. d. The light emitted by an electron moving from a higher to a lower energy level has a __________ directly proportional to the energy change of the electron. ground state excited quantum frequency

45. transition 2. The three groups of lines in the hydrogen spectrum correspond to the _________ of electrons from higher energy levels to lower energy levels.

46. 5.3 F. Quantum Mechanics 1. In 1905, Albert Einstein successfully explained experimental data by proposing that light could be described as quanta of energy. a. The quanta behave as if they were particles. b. Light quanta are called ________. 2. In 1924, De Broglie developed an equation that predicts that all moving objects have _________ behavior. photons wavelike

47. electron microscope 3. Today, the wavelike properties of beams of electrons are useful in magnifying objects. The electrons in an __________________ have much smaller wavelengths than visible light. This allows a much clearer enlarged image of a very small object, such as this mite.

48. 4. Classical mechanics adequately describes the motions of bodies much larger than atoms, while quantum mechanics describes the motions of __________ particles and atoms as waves. 5. The Heisenberg uncertainty principle states that it is __________ to know exactly both the velocity and the position of a particle at the same time. a. This limitation is critical in dealing with small particles such as electrons. b. This limitation does not matter for ordinary-sized object such as cars or airplanes. subatomic impossible

49. 7.3

50. Review Questions: Ans. (d) • Which of the following e- transitions is most likely to produce UV light in a hydrogen atom? (a) n = 3  n = 5 (b) n = 5  n = 3 (c) n = 1  n = 3 (d) n = 3  n = 1 2. Which series in the hydrogen atom produced the visible lines in the atomic spectrum? (a) Lyman (b) Brackett (c) Paschen (d) Balmer 3. Which color of visible light has the greatest amount of energy? (a) red (b) violet (c) green (d) yellow 4. Using the energy level diagram in your notes, determine which electron transition would result in the red line seen in the hydrogen spectrum. (a) n = 3  n = 2 (b) n = 4  n = 2 (c) n = 5  n = 2 (d) n = 6  n = 2 Ans. (d) Ans. (b) Ans. (a): Smallest gap btw levels