Chapter 5 Arrangement of electrons in atoms

1. Chapter 5Arrangement of electrons in atoms

2. *Rutherford's model of the atom does not explain how the electrons fill the space • Light (electromagnetic radiation) has a dual nature, meaning it behaves like a wave and a particle.

3. A) Wave description of light – • 1800's scientists believed that light was a beam of energy moving through space in the form of waves • (like waves on a lake when a pebble is thrown in)

4. *All waves have 4 characteristics: amplitude- height of wave origin to crest wavelength (λ)- distance between crests frequency (v) how fast up and down (oscillations) units: waves/sec, Hertz (Hz), s-1 speed (c) - constant 2.998 x 108 m/s

5. C = speed of light (latincelerata) Formula: c = λv* *wavelength and frequency are inversely proportional, meaning that if wavelength decreases then frequency increases & vice versa.

6. #1 Example problem: What is the frequency of light that has a wavelength of 450 nm? hint: convert nm to m (1m = 1 x 109 nm)

7. #2 Example problem: What is the wavelength of electromagnetic radiation if its frequency is 4.5 x 10-3 Hz?

8. Exit Question: 5 points • Write down 3 things that you learned today. • Write down one thing you don’t understand.

9. B) Particle description of light • 1900's experiments showed that light behaved like a stream of extremely tiny, fast moving particles.

10. 1) photoelectric effect - refers to the emission of electrons from a metal when light shines on the metal (but only if the frequency was at a certain minimum) ex/ solar powered items work if you have enough light

11. 2) Max Planck - studied light emitted from hot metal objects • (like a hot horseshoe glows). • He suggests that objects emit energy in small specific amounts called quanta.

12. 3) Quantum - minimum quantity of energy that can be lost or gained by an atom To calculate the energy of a quantum of light use formula: E = hv Where: E = energy (in Joules units) h = 6.626 x 10-34Js (Joule seconds) Planck's constant v = frequency

13. Albert Einstein (1905) • Introduces the wave-particle dual nature of light. • wave & particle behavior • each particle carries a quantum of energy. • EM radiation is absorbed by matter in whole numbers of photons.

14. photon - particle of light (EM radiation) having zero mass and carrying a quantum of energy. Ephoton = hv

15. Example problem: Using: Ephoton = hv Calculate the frequency for a photon of light that has an energy 3.2 x 10-19 J.

16. Hydrogen’s line emission spectrum • Niels Bohr passed electric current through hydrogen gas • PINK colored light emitted • When energy is added to an atom, electrons become excited & move to higher energy level.

17. A photon is emitted when the electrons move back to a more stable, GROUND state. • Ground state – lowest energy state of an atom • Excited state – state in which an atom has a higher potential energy than its ground state. • Ephoton= E2 – E1 = hv

18. Ephoton= E2 – E1 = hv The energy of this photon is equal to the difference in energy between the atom’s initial state and its final state.

19. Bohr Model of the Hydrogen atom 1913 • Bohr links the photon emission of hydrogen to a model of the atom’s electron. See p. 129 • Electron circles in orbits (defined paths) • Electron has a fixed energy • Each concentric circle orbit had an empty space in between where the electron could not exist (ladder analogy p. 129)

20. Explanation of the spectral lines produced by hydrogen: An electron cannot gain or lose energy. It can move to a higher energy orbit by gaining an amount of energy equal to the difference in final and initial states.

21. The Quantum Model of the atom • Quantum Theory – • modern description, • primarily mathematical, • of the behavior of electrons in atoms. • (it estimates the probability of finding an electron in a certain position)

22. Louis De Broglie (“de broylee”) 1924 He proposed an equation that suggested that any matter with mass and velocity has a corresponding wavelength.

23. Setting both energy equations equal to each other: • E = mc2 E=hv • mc2 = hv (substitute v with wavelength from c = λv) • Wavelength(λ) = h/mc

24. Werner Heisenberg 1927 • e- s are detected by their interaction with photons. • This interaction will change both the direction and position of the e-.

25. Heisenberg uncertainty principle • it is impossible to determine simultaneously both position and velocity of an e-

26. Therefore, e- s are located in orbitals or 3-D clouds of probable location • (not neat orbits like Bohr’s model)

27. Erwin Schrodinger came up with an equation that treated electrons in atoms as waves. • Quantization of electron energies was an outcome of his equation (vs. Bohr’s theory that assumed quantization as a fact) • Quantum numbers - numbers used to specify the energy, location, shape, orientation of atomic orbitals, & spins of electrons in orbitals

28. Quantum numbers - • Numbers used to describe an e- • 1. energy level • 2. location • 3. Shape of orbital • 4. orientation of orbitals • 5. spins of e-s in orbitals

29. 1) Principle quantum number– (n) main energy level occupied by the electron; the distance of an orbital to the nucleus. ex/ n= 1, n=2 (whole numbers) • 2) Angular momentum quantum number (l) indicates shape of orbital: s – sphere shape (2 e-) p – dumbbell shape (6e-) d – double dumbbell (10 e-) f – complex shape (14 e-)

30. 3) Magnetic quantum number (m) orientation of orbital (x, y, z) • 4) Spin Quantum number has two possible values +1/2 or -1/2

31. Electron Configurations - the ways in which electrons are arranged around the nucleus of an atom. Apply three rules: • Aufbau principle – electrons enter orbital of lowest energy first • Pauli exclusion principle – two electrons of opposite spin occupy an orbital (no two electrons in the atom can have the same set of quantum numbers)

32. 3) Hund’s Rule – When electrons occupy orbitals of equal energy, electrons fill the orbitals one at a time and then will pair up. Write the electron configuration for the following elements: • H b) B c) C d) Fe e) Zn Draw orbital diagrams for the same elements above.

33. Sec 1 • For electromagnetic radiation, c (speed of light) equals _________________________. • A quantum of electromagnetic energy is called _______________. • The energy of a photon is related to its _____________. • If electrons in an atom have the lowest possible energies, the atom is in the ________________. • Bohr’s theory helped explain why excited hydrogen gas gives off certain ___________ of light. • According to Bohr’s theory, an excited atom would _______________ energy.

34. Section 2 Review Q’s • A three-dimensional region around a nucleus where an electron may be found is called a(n) ____________. • Unlike in an orbit, in an orbital an electron’s position cannot be known _______________. • What are the 4 quantum numbers and what do they represent? • What are the shapes of the orbitals? • How many electrons fit in each orbital? • What is the difference between a 2s orbital and a 4s orbital?

35. Sec 2 • How many orbital shapes are possible at the 2nd energy level? 3rd energy level? • An electron for which n= 5 has more _____ than an electron for which n=3. • If 8 electrons completely fill a main energy level, what is n?

36. Section 3 Review Q’s • Draw the diagonal rule. What does this rule show? • Know the 3 rules for writing electron configurtions. • Write the electron configuration for Si. • Draw the orbital diagram for Mg. • What element has the following configuration: 1s22s22p63s1 ? • How many electrons in the highest energy level of a bromine atom? • Which element has the electron configuration of [Ar]4s23d104p5