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-history continued. Rutherford’s model of the atom slightly wrong-why? Electrons and protons are attracted to each other. In order for the electrons to stay away from what they are attracted to, they need energy. Neils Bohr introduced a new model that had the electrons located on energy levels. Electrons filled the energy level closest to the nucleus first as it was the cheapest to fill in terms of energy. Then the electrons filled then next energy level, and so on. Each energy level could only hold a certain number of electrons.

4. History of the Atom Continued • Bohr’s Planetary Model-7 energy levels • Drawing • -first energy level can hold 2 electrons. • -second energy level can hold 8 electrons. • -third energy level can hold 18 electrons. • -fourth-seventh energy level could hold 32 electrons. Example of how to draw • Lithium-7 Lithium has 3 electrons-2 on first energy level 1 and 1 on the second. Protons and neutrons in the middle.

5. History of the Atom Continued • Bohr’s Planetary Model-7 energy levels • Drawing • -first energy level can hold 2 electrons. • -second energy level can hold 8 electrons. • -third energy level can hold 18 electrons. • -fourth-seventh energy level could hold 32 electrons. 2nd example –You try it!! • Oxygen-16 oxygen has 3 electrons-2 on first energy level 1 and 6 on the second. Protons and neutrons in the middle. • Quantum- the energy needed for electrons to move between energy levels. Add energy to move away from nucleus and release energy to move toward the nucleus.

6. Quantum Mechanical Model • Bohrs model only worked for elements 1-18 and then something happened. Potassium-39 had 19 electrons did not follow the model. • There are 2 electrons on the first energy level. There are 8 on the second, 8 on the third, and 1 on the fourth. It did not fill up the third with 18. Why did it put the last electron in the 4th level? This model must be wrong..

7. New Model: 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.

8. Quantum Number: Like an address • Principle Energy Level: n : “Like a __________.”

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

10. 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

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

12. Rules for Electron Configuration

13. 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 ...

14. 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

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

16. 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.

17. 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

18. Light and the Atomic Spectra Crest Amplitude Wavelength Trough

19. 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

20. 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

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

22. Example 1a

23. Example 1b

24. 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.

25. 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

26. Example 2a

27. Example 2b

28. Energy Traveling Towards Nucleus

29. 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.

30. Chapter 6 The Periodic Table

31. History of the Periodic Table

32. 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.

33. 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.

34. Mendeleev’s Table

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

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

37. Classification Systems for the Periodic Table

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

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

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

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

42. 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”

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

44. Groups/Families And their Properties

45. 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.

46. 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.

47. 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.

48. 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.

49. 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

50. 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