Unit 3

1. Unit 3 Electrons & The Periodic Table Chapter 5 & 6

2. Chapter 5 Electrons in Atoms

3. History of the Atom Continued • Dalton, Thomson, Rutherford • Bohr’s Planetary Model • Quantum- the energy needed to move between energy levels.

4. Quantum Mechanical Model • The address on an electron based on the Schrodinger Probability Equation. • “The likely hood of finding an electron’s path (location).” • Heisenberg Uncertainty Principle • It is impossible to know both the velocity and position of a particle at the same time.

5. Quantum Number: Address • Principle Energy Level: n : “Like a __________.”

6. Quantum Number Continued • Sublevel • “Like a ________” • Levels within a Principle Energy Level that include_______________. • Orbitals: paths that electrons follow within a sublevel • “Like a _________.”

7. Atomic Orbitals • A region of space around the nucleus of an atom where there is a high probability of finding an electron. • Sublevel Orbital drawing # electrons • s • p • d • f

8. Electron Configuration • The arrangement of electrons around the nucleus of an atom in its ground state. • The address of an electron.

9. Rules for Electron Configuration

10. The Aufbau Principle: Rule-1 • Lowest energy level first. • The energy ordering of the orbitals can be remembered from this diagram: If you follow this line through the diagram, it traces out the sub shells in this order:  1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p ...

11. 1s Least Energy 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f Most Energy

12. Pauli Exclusion Principle: Rule 2 • No more than two electrons can occupy an atomic orbital.

13. Hund’s Rule: Rule-3 • When electrons occupy orbitals of equal energy, one electron occupies each orbital until all orbitals contain one electron with their spins parallel. (all up or down) • Electrons separate out so that each orbital at the same energy level has one electron before the electrons start to double up.

14. Exceptions to the rules: • Move electrons to have full or ½ full orbitals; more stable configuration this way • d4 = Take an electron from the previous s orbital and make the s and d orbitals 1/2 full. • Becomes _____s1 ____d5 • d9 = Takes an electron from the previous s orbital and makes the s 1/2 full and the d full. • Becomes _____s1_____d10

15. Light and the Atomic Spectra Crest Amplitude Wavelength Trough

16. Light and the Atomic Spectra Cont. • Spectrum • The range of wavelengths and frequencies making up light. • Electromagnetic Radiation • A series of energy waves that travel in a vacuum at 2.998 x 108 m/s

17. Frequency • Frequency: • Number of wavelength per second • Units = 1/seconds or s-1 or hertz • Wavelength: length of one wave: crest to crest, trough to trough, or above. The units are usually in meters or nanometers. • The relationship between frequency and wavelength is that as one get bigger, the other gets smaller. “inverse relationship”. 1 second time interval

18. Equation 1: • Frequency x Wavelength = 2.998 x 108 (1/s) (m) (m/s)

19. Example 1a

20. Example 1b

21. Movement of Electrons • The energy produced by one packet of electrons called photons that are moving back towards the nucleus. • The electrons are releasing excess energy from being excited by heat, electricity, or pressure, as electromagnetic radiation. • Ground State • The lowest energy level occupied by an electron when an atom is in its most stable energy state. • Excited state • An electron in a higher energy level than the ground state.

22. Equation 2: • Frequency x 6.626 x 10-34 = Energy (1/s) (J*s) (J) • Plank’s Constant • A number used to calculate the radiant energy absorbed or emitted by a body from the frequency of radiation. • 6.626 x 10-34 J*s

23. Example 2a

24. Example 2b

25. Energy Traveling Towards Nucleus

26. Atomic Emission Spectra • The pattern of frequencies obtained by passing light emitted by atoms of elements in the gaseous state through a prism. • Using a spectroscope (prism) to view electromagnetic radiation produced by excited atoms. • A fingerprint of the atom is seen and used for identification.

27. Chapter 6 The Periodic Table

28. History of the Periodic Table

29. Dimitri Mendeleev (1834-1907) • Russian chemist • First to arrange the elements in a logical way. • Arranged elements in order of increasing atomic mass. • Grouped elements in families according to chemical and physical properties.

30. Mendeleev continued • Left blank spaces in the table where there were no known elements with the appropriate properties and masses. • Predicted physical and chemical properties of the missing elements.

31. Mendeleev’s Table

32. Henry Moseley (1887-1915) • British physicist • Rearranged elements according to atomic mass instead of atomic mass. • Fixed the errors in Mendeleev’s table.

33. Periodic Law • When the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

34. Classification Systems for the Periodic Table

35. In the Beginning • Elements were classified by • metals • nonmetals • metalloids

36. Properties of Metals • high electric conductivity • high luster when cleaned • ductile • Can be drawn into wires • malleable • Can be hammered into thin sheets

37. Properties of Nonmetals • Nonlusterous - Not shinny. • Poor conductors of electricity • Can be gases at room temperature • Some are brittle solids

38. Properties of Metalloids • Properties that are intermediate between those of metals and nonmetals

39. After the arrangement of the periodic table: • Elements were classified as • Representative Elements – Group A elements • Transition Elements – Group B elements • Inner Transition Elements – last two “rows”

40. Organization of the Periodic Table • Periods – Horizontal rows • Groups/families – share similar properties; vertical columns in the periodic table

41. Groups/Families And their Properties

42. Alkali Metals • In Group IA of the periodic table. • The members of the family • lithium, sodium, potassium, rubidium, cesium, and francium. • All six elements have the properties of metals except they are softer and less dense. • Can be cut with a knife. • Most reactive metals. • So reactive that they are never found in nature. • Always combined with other elements.

43. Alkaline Earth Metals • In Group 2A of the periodic table • Members are: • beryllium, magnesium, calcium, strontium, barium, and radium. • Harder and more dense than the alkali metals • Have higher melting points and boiling points. • Highly reactive, but not as active as the alkali metals. • Never found free in nature.

44. Halogens • In family VIIA. • They are strongly nonmetallic. • Members are • fluorine, chlorine, bromine, iodine, and astatine. • They are the most active nonmetals. • They have low melting points and boiling points. • In the gas phase, they exist as diatomic elements. • Halogens combine readily with metals to form a class of compounds known as salts.

45. Noble Gases • In group 8A. • Colorless gasses that are extremely unreactive. • Do not readily combine with other elements to form compounds, the noble gasses are called inert. • The family of noble gasses includes • helium, neon, argon, krypton, xenon, and radon. • All the noble gasses are found in small amounts in the earth's atomsphere. • One important property of the noble gasses is their inactivity.

46. Transition Metals • Most transition metals are excellent conductors of heat and electricity. • Most have high melting points and are hard. • Transition metals are much less active than the alkali and alkaline earth metals. • Many transition metals combine chemically with oxygen to form compounds called oxides. • Transition metals form compounds that are brightly colored

47. Inner Transition Metals • Have similar properties to transition metals • Are located below the periodic table. • Members in this group have atomic numbers 57-76 & 89-102

48. Electron Configurations and the Periodic Table • Periods • energy level for s and p, d is shifted 1, f is shifted 2 • Groups/Families • 1A : s1 • 2A : s2 • 3A : p1 • 4A : p2 • 5A : p3 • 6A: p4 • 7A : p5 • 8A : p6 • Transitions : d’s • Inner transition/rare earth : f’s

49. Other Periodic Trends

50. Atomic Size • ½ the distance between the nuclei of two atoms of an element. • Row • Increases from right to left • Family/Group • Increases as you move down a column