1. Chapter 13 Electrons in Atoms C. Smith

2. I. Models of the Atom A. The Evolution of Atomic Models 1. There are four major models of the atom that have been developed from John Dalton theory. 2. Dalton Atomic Theory a. He theorized that an atom was indivisible, uniformly dense sphere. b. He theorized that all atoms of the same element have the same mass and the same chemical behaviors. c. He theorized that atoms of different elements have different chemical behaviors. d. He theorized that atoms of different elements combine to form compounds. (Example — H 2 O)

3. I. Models of the Atom A. The Evolution of Atomic Models 3. J.J. Thomson realized that the accepted model did not take electrons into account. a. He is credited with the discovery of the negatively charged particles called electrons. b. He theorized that the atom is a dense sphere with a positive charge and also contains negative charged particles. c. His model is also known as the Plum Pudding model.

4. I. Models of the Atom A. The Evolution of Atomic Models 4. Ernest Rutherford expanded on Thomson’s theory. a. The atom has a very dense center of positive charge called the nucleus. b. The nucleus contains the protons for the atom and make up more than 99.9% of its mass. c. The electrons surround the nucleus.

5. I. Models of the Atom A. The Evolution of Atomic Models 5. Niels Bohr proposed a model in which the electrons move around the nucleus. a. He theorized that the electron orbits the nucleus. b. He theorized that the orbits were different energy levels that the electrons travel in and can be excited to a high energy level. c. He theorized that the electrons did not lose energy and fall into the nucleus. (The weakness in Rutherford’s theory.)

6. I. Models of the Atom A. The Evolution of Atomic Models 6. A quantum of energy is the amount of energy required to move an electron from its present energy level to the next higher one. (Also called a quantum leap)

7. I. Models of the Atom B. The Quantum Mechanical Model 1. Erwin Schrödinger related the amplitude of the electron wave,  (psi), to any point in space around the nucleus. 2. His equation treated the electron as a wave and developed an equation to describe this behavior.

8. I. Models of the Atom B. The Quantum Mechanical Model 3. The quantum mechanical model comes from the mathematical solutions to Schrödinger equation. 4. The quantum mechanical model does not define an exact path for the electron to take around the nucleus but instead estimates a probability of finding the electron in a certain position. 5. Since the volume occupied by an electron is somewhat vague, it is better to refer to an electron cloud.

9. I. Models of the Atom B. The Quantum Mechanical Model

10. I. Models of the Atom C. Atomic Orbitals 1. Electrons can occupy only specific energy levels. 2. These energy levels, referred to as “n” is called the principal quantum number. 3. The maximum number of electrons that a level can contain is 2n 2 (Whole number integers only).

11. I. Models of the Atom C. Atomic Orbitals 4. These are referred to as sublevels and the number of sublevels for each energy level is equal to the value of the principal quantum number. 5. The lowest energy level is “s”. 6. The second lowest is “p”,the third lowest level is “d”, and the remain level is “f”.

12. I. Models of the Atom C. Atomic Orbitals 7. The “s” orbital is spherical in shape and contains 2 electrons and is also called the ground state. 8. The “p” level is barbell shape and exist along the axis of the plane. 9. The “d” orbitals exist in the plane. 10. The “s” level contains 1 pair of electrons, “p” contains 3 pairs, “d“ contains 5 pairs, and “f” contains 7 pairs.

13. II. Electron Arrangement in Atoms A. Electron Configurations 1. The ways in which electrons are arranged around the nucleus is called electron configuration. 2. The are three rule that tell you how to find the configurations. a. Aufbau principle b. Pauli Exclusion principle c. Hund’s Rule

14. II. Electron Arrangement in Atoms A. Electron Configurations 2a This is called the Aufbau principle. 1. Electrons enter at the lowest energy level. 2. Some energy levels overlap into the adjacent principal energy level.

15. II. Electron Arrangement in Atoms A. Electron Configurations

16. 2b. This is called the Pauli exclusion principle. 1. Spectral data shows that only 2 electrons can exist in the same orbital. 2. Electrons behave as if they were spinning about their own axis. 3. When electrons occupy the same orbital – they are said to spin in opposite directions (assign +1/2 and – 1/2).

17. II. Electron Arrangement in Atoms A. Electron Configurations 2c. This is called Hund’s Rule. 1. Also with the principle, you must have all orbital filled with one electron before you can add the other electron with opposite spin to the orbital. 2. All elements would like to have a completely filled orbital and the maximum number of electrons that can exist in a filled orbital is eight.

18. II. Electron Arrangement in Atoms A. Electron Configurations 3. When writing electron configurations, you must know the total number of electrons for the element (atomic number). 4. Write down the sequence of orbitals. 5. Draw circle to represent the orbitals. 6. Place arrows (or slashes) to represent the electrons.

19. II. Electron Arrangement in Atoms A. Electron Configurations

21. II. Electron Arrangement in Atoms B. Exceptional Electron Configurations 1. Filled sublevels are more stable than partial filled or half-filled sublevels. 2. But sometimes half-filled may be more stable than other configurations.

22. III. Physic and the Quantum Mechanical Model A. Light and Atomic Spectra 1. This energy consist of variation in electric and magnetic fields taking place in a regular, repeating fashion. (Electromagnetic energy) 2. Light is a form of electromagnetic radiation

23. III. Physic and the Quantum Mechanical Model A. Light and Atomic Spectra 3. If you plot the strength of the variation against time, the graph shows “waves” of energy. 4. The number of waves peaks that occur in a unit of time is called the frequency of the wave (Greek letter v and units are Hertz (Hz)).

24. III. Physic and the Quantum Mechanical Model A. Light and Atomic Spectra 5. The distance between the peaks is the wavelength (Greek letter λ) and the amplitude of a wave is the height from the maximum displacement from zero. 6. These characteristics of waves are related by the statement c= λv where c is the speed of light which is 3.0 x 10 8 m/s.

25. III. Physic and the Quantum Mechanical Model A. Light and Atomic Spectra 7. The wavelengths of light can separate into a spectrum of colors. 8. This is part of the visible spectrum. 9. There are two types of spectrums. a. Adsorption spectrum. b. Emission spectrum.

26. III. Physic and the Quantum Mechanical Model A. Light and Atomic Spectra 10. Adsorption spectrum is when the energy gained by the excited electron is is absorbed so that it is missing in visible spectrum. 11. Emission spectrum is when the excited electrons lose the energy and it is emitted at specific points on the visible spectrum that appear as single lines on a detector.

27. III. Physic and the Quantum Mechanical Model B. The Quantum Concept and the Photoelectric Effect 1. Max Planck used Bohr’s theory to develop his hypothesis. 2. He assumed that energy is given off in packets called quanta or photons instead of a steady stream.

28. III. Physic and the Quantum Mechanical Model B. The Quantum Concept and the Photoelectric Effect 3. He stated that the amount of energy given off is related to the frequency of light (v - Greek letter nu). 4. He thought a quantum energy was equal to E = hv where h is the constant 6.63 x 10 -34 J/Hz (Hz = Hertz).

29. III. Physic and the Quantum Mechanical Model B. The Quantum Concept and the Photoelectric Effect 5. Albert Einstein proposed that light could be described as a quanta of energy that behaved as if they were particles. 6. The dual wave-particle behavior is called the photoelectric effect.

30. III. Physic and the Quantum Mechanical Model B. The Quantum Concept and the Photoelectric Effect 7. In the photoelectric effect, metals eject electrons when light shines on them. 8. The frequency and the wavelength of the light determine if the photoelectric effect will occur.

31. III. Physic and the Quantum Mechanical Model C. An Explanation of Atomic Spectra 1. Consider the electron of a hydrogen atom in its lowest energy level, or ground state. 2. The quantum numbers represent the different energy states.

32. III. Physic and the Quantum Mechanical Model C. An Explanation of Atomic Spectra 3. The difference between these energy states corresponds to the lines in the hydrogen spectrum. 4. With more complex atoms more than one electron is present and the interaction between electrons make solution to the equation impossible because electrons have the same charge.

33. III. Physic and the Quantum Mechanical Model C. An Explanation of Atomic Spectra 5. It is possible to approximate the electronic structure of a multi-electron atom. 6. This approximation is made by first calculating the various energy states and quantum numbers. 7. It is assumed that the various electrons in multi-electron atom occupy the same energy states without affecting each other.

34. III. Physic and the Quantum Mechanical Model D. Quantum Mechanic 1. Louis De Broglie proposed an idea based on Planck’s theory and Einstein’s relationship of matter and energy. 2. Using the two formulas, he equated mc 2 = hν (v = frequency).

35. III. Physic and the Quantum Mechanical Model D. Quantum Mechanic 3. He substituted v * for the velocity of light (c) so that mv * 2 = hv and v /λ for v to get mv * 2 = hv /λ. (λ - Lamda = wavelength) 4. To determine wavelength (λ), the equation becomes λ = h/mv.

36. III. Physic and the Quantum Mechanical Model D. Quantum Mechanic 5. This allows for predictions of the wavelength of a particles. 6. Werner Heisenburg refined ideas about atomic structure. 7. He stated that it is impossible to know the exact position and momentum of an electron in an atom.

37. III. Physic and the Quantum Mechanical Model D. Quantum Mechanic 8. Using the equation for momentum, he proposed that mv = p where m is mass and p is momentum. 9. The uncertainty of position and momentum are related to Planck’s constant ∆p ∆x > h where p is momentum and x is position (∆ = change). 10. Because h is constant, ∆ p and ∆ x are inversely proportional to each other.