1. CHAPTER 4 CHEMISTRY

2. PROPERTIES OF LIGHT (P91-93) Originally thought to be a wave It is one type of ELECTROMAGNETIC RADIATION (exhibits wavelike behavior as it travels through space) Ex. X-ray, ultraviolet and infrared light, microwaves, and radio waves

3. FEATURES OF WAVES Wavelength Distance between corresponding points on adjacent waves Frequency (v) The number of waves that pass a point in a specific time (usually one second)

4. PHOTOELECTRIC EFFECT (P93-94) Emission of electrons from a metal when light shines on the metal Because the light is needed to be a certain frequency before it can emit an electron, the wave theory didn’t fit (like a soda machine that takes quarters only) Max Planck suggested that light is emitted in small specific amounts called QUANTA

5. MAX PLANCK Quantum is the minimum quantity of energy that can be lost or gained by an atom E = hv E = energy in joules v = frequency h = Planck’s constant: 6.626 x 10 -34 J s

6. ALBERT EINSTEIN Said that electromagnetic radiation has a dual wave- particle nature Each particle of light is called a photon Energy of light depends on frequency E photon = hv

7. PRACTICE PROBLEMS Wavelength (m)Frequency (Hz)Energy (J).001 7.0 x 10 13 5.0 x 10 -7 2.0 x 10 -15 1.2 x 10 22 C = ƛ vE = hv C = speed of light = 3.0 x 10 8 m/s H = planck’s constant = 6.626 x 10 -34 J s

8. ANSWERS Wavelength (m)Frequency (Hz)Energy (J).0013.0 x 10 11 2.0 x 10 -22 4.3 x 10 -6 7.0 x 10 13 4.6 x 10 -20 5.0 x 10 -7 6.0 x 10 14 4.0 x 10 -19 1.0 x 10 -10 3.0 x 10 18 2.0 x 10 -15 2.5 x 10 -14 1.2 x 10 22 8.0 x 10 -12 C = ƛ vE = hv C = speed of light = 3.0 x 10 8 m/s H = planck’s constant = 6.626 x 10 -34 J s

9. NOTICE THE PATTERNS Wavelength (m)Frequency (Hz)Energy (J).0013.0 x 10 11 2.0 x 10 -22 4.3 x 10 -6 7.0 x 10 13 4.6 x 10 -20 5.0 x 10 -7 1.6 x 10 14 1.1 x 10 -18 1.0 x 10 -10 3.0 x 10 18 2.0 x 10 -15 2.5 x 10 -14 1.2 x 10 22 8.0 x 10 -12 What do you notice about wavelength? What do you notice about frequency? What do you notice about energy? What happens to wavelength if frequency goes up? What happens to energy if frequency goes up?

10. HYDROGEN ATOM LINE EMISSION (P.94,95) Ground state = the lowest energy state of an atom Excited state = an atom has a higher potential energy than it has in its ground state Current passes through gas (like neon) and excites the gas; when it returns to its ground state it gives off electromagnetic radiation (like the colors of neon light)

11. HYDROGEN ATOM LINE EMISSION When a current is passed through hydrogen gas it gave off a pink glow. Line emission spectrum = the different wavelengths of the pink glow when passed through a prism This told scientists that each atom has specific energy levels that electrons jump to and from when going from excited state to ground state

12. HYDROGEN ATOM LINE EMISSION Every element has a characteristic spectrum Spectroscopy = identifying a substance by the wavelengths that are separated by a prism

13. BOHR’S MODEL (P. 96-97) Earlier models: Cubic model Plum pudding model (thomson’s) Saturnian model Rutherford’s model

14. 1913 (Niels Bohr) proposed a new model that explained the hydrogen atom line emission spectrum Also called planetary model Electrons orbit nucleus at different levels Closer to the nucleus means lower energy level; farther away is higher energy level

15. PROBLEMS WITH BOHR’S MODEL This model could NOT explain the spectra of atoms with more than one electron It did not fully explain the chemical behavior of atoms

16. QUANTUM MODEL Electrons behave like waves (De Broglie) Electrons are confined to the space around an atomic nucleus (like the frequency of wave) They can be bent (or diffracted) They can interfere with each other (overlapping and reduction: p.99 fig.4-10)

17. SCHRODINGER WAVE EQUATION Erwin Schrodinger used the dual wave-particle theory to develop an equation that treated electrons in atoms as waves. This with the Heisenberg Uncertainty principle laid the foundation for modern quantum theory

18. QUANTUM THEORY Describes mathematically the wave properties of electrons and other very small particles The wave function (schrodinger’s equation) gives only the probability of finding an electron at a given place around the nucleus

19. ORBITALS Bohr’s idea about neat orbit’s around the nucleus was wrong An orbital is a 3 dimensional region around the nucleus that indicates the probable location of an electron

20. HEISENBERG UNCERTAINTY It is impossible to determine simultaneously both the position and velocity of an electron

21. ATOMIC ORBITALS AND QUANTUM NUMBERS Quantum numbers specify the properties of atomic orbitals and the properties of electrons in orbitals. There are four quantum numbers:

22. PRINCIPAL QUANTUM NUMBER Symbolized by n Tells the main energy level occupied by the electron Only positive numbers as n increases, energy increases

23. AS N INCREASES, ENERGY INCREASES

24. ANGULAR MOMENTUM QUANTUM NUMBER Symbolized by l Indicates the shape of the orbital Shapes are indicated by Sspherical Pdumbell shaped

25. ANGULAR MOMENTUM QUANTUM NUMBER (CONTINUED) D pair of dumbells mostly (p. 103)

26. ANGULAR MOMENTUM QUANTUM NUMBER (CONTINUED) F very complicated

27. MAGNETIC QUANTUM NUMBER Symbolized by m Indicates the orientation of an orbital around the nucleus In S level m=0 In P orbital m= -1, 0, or +1 In D; m=-2, -1, 0, 1, or 2 In F: there are 7 possible orientations

28. SPIN QUANTUM NUMBER Only 2 possible values (+1/2 or -1/2) One orbital can hold 2 electrons but they must have the opposite spin

29. ELECTRON CONFIGURATION The arrangement of electrons in an atom 3 basic rules

30. AUFBAU PRINCIPLE An electron occupies the lowest energy orbital that can receive it

31. PAULI’S EXCLUSION PRINCIPLE No two electrons in the same atom can have the same set of four quantum numbers

32. HUND’S RULE Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin