Chapter 5 Review

1. Chapter 5 Review

2. Wave Nature of Light • Wavelength- The distance between two consecutive peaks or troughs. • Frequency- The number of waves that pass a point each second (the unit is the Hertz, Hz). One complete wave or cycle per. second = 1 Hz. • Velocity- Distance a peak moves in a unit of time.

4. Particle Nature of Light PHOTOELECTRIC EFFECT • When light shines on certain metals an electric current is produced (electrons are emitted from the metal). • Not explained by the wave theory of light… • Light is thought of as a stream of particles called…

5. Bohr’s Model • Atom absorbs energy  electron ‘jumps’ to a higher energy level (orbit). • Electron ‘jumps’ to a lower orbit  energy released (photon). • Energy difference between the two orbits  energy of the photon The color (frequency) of light produced E = h f

6. Bohr’s Shortcomings • Quantized energy levels were not immediately accepted by other scientists (though correct) • Worked for hydrogen but not for other elements with more electrons.

7. Heisenberg Uncertainty Principle • It is impossible to determine simultaneously the position and velocity of an electron or any other particle.

8. Orbital • A three dimensional region around the nucleus that indicates the most probable location of the electron.

9. Quantum Numbers • Describe the properties of atomic orbitals and the electrons that occupy them… Quantum NumberDescribes • Principle - main E. level • Angular Momentum - shape of Orbital • Magnetic - orientation of orbital • Spin - orientation of electron

10. Rules for Writing e- configurations • An electron occupies the lowest energy orbital that can receive it…(like filling a glass) • If two electrons occupy the same orbital they must have opposite spins! • Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron.

11. 1S2 2S2 2P2

12. What do the symbols represent?

13. Write Orbital notation for • Nitrogen

14. Electron configuration for: • Ne 1S2 2S22P6

15. Nobel Gas Notation for: • Aluminum • Al [Ne] 3S2 3P1

16. How does the arrangement of electrons relate to an elements position on the periodic table? • ….where would find an element with an outermost energy level like this 4S24P4 ?

17. Calculate Frequency and energy • Wavelength = 4.5 x 10-7 meters

18. Region outside of nucleus where an electron can most probably be found • Orbital

19. An electron in the lowest possible energy level said to be in it’s • Ground state

20. When an electron absorbs energy it will be in… • An excited state

21. A photon is released when an electron moves from • Excited state to ground state

22. A quantum of electromagnetic energy is a • Photon

23. How does an orbit differ from an orbital? • Orbitals do not show the exact location or path of an electron… only the probable location

24. It is not possible to know the precise location and velocity of an electron or other small particle (“what is…”) • Heisenberg uncertainty principle

25. Indicates the main energy level of an orbital or electron • Principle quantum number

26. An electron at main energy level 5 has more ___ than an electron at main energy level 2 • Energy

27. Dumbell-shaped set of 3 orbitals • P ORBITALS

28. Indicates the orientation of the orbital with regard to a three dimensional axis • Magnetic quantum number

29. A spherical electron cloud • S orbital

30. The difference between a 2S orbital and a 3S orbital • Distance from nucleus or energy level

31. An electron occupies the lowest energy orbital that can receive it • Aufbau principle

32. Orbitals of equal energy are each occupied by one electron before any is occupied by a second electron. • Hund’s Rule

33. Two electrons in the same orbital must have opposite spins • Pauli exclusion principle

34. The Bohr model was an attempt to explain what about hydrogen • Bright line or line emission spectrum

35. The number of electrons in the highest occupied energy level of argon (a noble gas) • 8

36. The distance between successive peaks on a wave • wavelength