Electrons in Atoms

1. Electrons in Atoms

2. Models of the Atom – A History • John Dalton • atom was solid, indivisible mass • J.J. Thomson • “plum pudding” model • e- stuck in lump of + charged matter • Ernest Rutherford • discovered nucleus • lacked detail about how electrons occupy the space surrounding the nucleus • did not address why the charged electrons are not pulled into the atom’s nucleus

3. Niels Bohr • e- in circular paths around nucleus • “planetary model” • e- have fixed energy

4. In the early 1900s, scientists began to unravel the puzzle of chemical behavior. • They had observed that certain elements emitted visible light when heated in a flame. • Analysis of the emitted light revealed that an element’s chemical behavior is related to the arrangement of the electrons in its atoms. • So…we need to understand a little about light!

5. Light and Energy • light: viewed as a wave and a particle (Isaac Newton) • electromagnetic radiation: any kind of light, visible or not • amplitude: height of a wave • wavelength: distance from crest to crest • frequency: # of waves that pass a point in a given time • measured in Hertz (Hz)

6. Amplitude

7. All electromagnetic waves, including visible light, travel at a speed of 3.00 x 108 m/s in a vacuum. • The speed of light is the product of its wavelength (λ) and its frequency (ν). • Although the speed of all electromagnetic waves is the same, waves may have different wavelengths and frequencies

8. As you can see from the equation, wavelength and frequency are inversely related; in other words, as one quantity increases, the other decreases

9. Electromagnetic Spectrum

10. Particle Nature of Light • The wave model of light cannot explain why heated objects emit only certain frequencies of light at a given temperature, or why some metals emit electrons when colored light of a specific frequency shines on them. • Obviously, a totally new model or a revision of the current model of light was needed to address these phenomena.

11. The quantum concept • In 1900, the German physicist Max Planck (1858–1947) began searching for an explanation as he studied the light emitted from heated objects.

12. The quantum concept • His study of the phenomenon led him to a startling conclusion: matter can gain or lose energy only in small, specific amounts called quanta. • That is, a quantum is the minimum amount of energy that can be gained or lost by an atom. • Matter can have only certain amounts of energy—quantities of energy between these values do not exist.

13. Electrons and Light • ground state: an e- in its lowest energy state • excited state: an e- in a higher than normal energy level

14. energy level: region around nucleus where e- is likely to be moving • e- can move up or down but cannot exist between levels • e- must gain right amt. of E to move to a higher level (lose E to go down) BOHR MODEL ONLY HOLDS TRUE FOR THE HYDROGEN ATOM!!!!

16. Atomic Emission Spectra • The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by atoms of the element.

17. An atomic emission spectrum is characteristic of the element being examined and can be used to identify that element. • The fact that only certain colors appear in an element’s atomic emission spectrum means that only certain specific frequencies of light are emitted.

18. Emission Spectrum for Hydrogen • In your reference tables • Only ABSORB energy when e- moves to an excited state • E is EMITTED if an e- moves to a lower level

19. Emission Spectrum for Hydrogen

21. Emission Spectra for H, Hg, Ne

22. Quantum Mechanical Model • Erwin Schrödinger • complex mathematical formula • no definite path for e- • only gives probability of finding e- • Heisenberg Uncertainty Principle: it is impossible to know the exact location and speed of an electron at any time

23. Quantum Theory • Four quantum numbers exist – n, l, m, s I. Principal Quantum Number (n) • energy level number • indicates size of electron cloud • n = whole # > 0 • # of e- in each level = 2n2 • How many e- can exist in levels 1-5?

24. II. Sublevels (l) - Azimuthal Quantum # • l = 0 to n-1 • indicates shape of the e- cloud • sublevels called s, p, d, and f • energy level # tells # of sublevels • Ex. Energy level 1 (n=1), 1 sublevel Energy level 2 (n=2), 2 sublevels • How many sublevels exist in energy levels 3-5?

25. III. Orbitals (m) – Magnetic Quantum # • regions where e- are likely to be found • each orbital can hold 2 e- • each sublevel has its own specific orbitals • m = -l to +l • s sublevel = 1 orbital = 1 pair e- (2) • p sublevel = 3 orbitals = 3 pair e- (6) • d sublevel = 5 orbitals = 5 pair e- (10) • f sublevel = 7 orbitals = 7 pair e- (14)

26. Orbital Shapes • s – spherical

27. p – dumbbell shaped

28. d orbital – clover-leaf shaped

29. f orbitals are too complex to be visualized

30. IV. Spin (s) • s = +1/2 or -1/2 • each e- in orbital must spin in opposite direction – WHY? • one clockwise, one counterclockwise

31. Arrangement of Electrons • Electron Configurations • ways in which e- are arranged around the nucleus • high E is unstable • unstable systems lose E to become more stable

32. Aufbau Principle • e- enter orbitals of lowest E first • s is lowest E, f is highest E • Pauli Exclusion Principle • no 2 e- can have the same set of quantum numbers • Hund’s Rule • e- will occupy empty orbitals of equal E before pairing up in an orbital

34. Exceptional Electron Configurations • half-filled and filled sublevels are more stable than partially filled sublevels • e- will shift to become more stable • transition metals are affected • Examples (you need to know!) • Copper • Silver • Chromium