1. Atomic Structure

2. Early Theories of Matter Science as we know it did not exist several thousand years ago Most philosophers thought that matter was made up of air, water, earth, and fire

3. Democritus (460-370 BC) The first to propose that matter was not infinitely indivisible Believed that matter was made up of atomos (or atoms) He also said that atoms could not be created or destroyed

4. Democritus’ Theory 1.Matter is composed of empty space through which atoms move 2.Atoms are solid, homogeneous, indestructible, and indivisible 3.Different atoms have different sizes and shapes 4.The differing properties of matter are due to the size, shape, and movement of atoms 5.Changes in matter result from changes in the groupings of atoms and not the atoms themselves

5. Democritus Democritus was on the right track, but had no scientific data to back up any of his claims The foremost thinker of the day, Aristotle, rejected Democritus’ ideas because he believed that nothingness could not exist No other theories of the atom came about until nearly 2000 years later

6. John Dalton John Dalton was the next scientist to propose a theory about the atom in the 19 th century

7. Dalton’s Atomic Theory 1.All matter is composed of extremely small particles called atoms 2.All atoms of a given element are identical, having the same size, mass, and chemical properties. Atoms of a specific element are different from other elements 3.Atoms cannot be created, destroyed, or broken into smaller particles 4.Different atoms combine in simple whole number ratios to form compounds 5.In a chemical reaction, atoms are separated, combined, or rearranged

8. Dalton’s Atomic Theory This was a HUGE step toward the atomic model, but was not totally correct Atoms ARE divisible into smaller subatomic particles Atoms of a given element can have different masses

9. JJ Thomson JJ Thomson used the cathode ray experiment to determine the charge to mass ratio of an electron. He identified the first subatomic particle, the electron He proposed the plum pudding model of the atom Credited for discovering the electron

10. Robert Millikan Millikan is noted for his famous Millikan’s Oil Drop Experiment This experiment determined the charge and the mass of an electron

11. Millikan's Oil Drop Experiment Click in this box to enter notes. Copyright © Houghton Mifflin Company. All rights reserved. Go to Slide Show View (press F5) to play the video or animation. (To exit, press Esc.) This media requires PowerPoint® 2000 (or newer) and the Macromedia Flash Player (7 or higher). [To delete this message, click inside the box, click the border of the box, and then press delete.]

12. Earnest Rutherford Rutherford’s Gold Foil Experiment helped to determine the existence of the nucleus Rutherford proposed that the nucleus was small, dense and positively charged Proposed the nuclear atomic model which stated that there was a nucleus with a positive charge and electrons around the outside

13. Gold Foil Experiment Click in this box to enter notes. Copyright © Houghton Mifflin Company. All rights reserved. Go to Slide Show View (press F5) to play the video or animation. (To exit, press Esc.) This media requires PowerPoint® 2000 (or newer) and the Macromedia Flash Player (7 or higher). [To delete this message, click inside the box, click the border of the box, and then press delete.]

14. James Chadwick Chadwick showed that the nucleus also contained neutrons He is credited for the discovery of the neutron

15. Atomic Structure

16. Basic Definitions Atom – smallest unit of an element that retains the properties of that element Atoms are made up of several subatomic particles called protons, neutrons, and electrons

17. Protons, Neutrons, & Electrons Protons – have a +1 charge and are found in the nucleus of the atom Neutrons – have no charge and are also found in the nucleus of an atom Electrons – have a -1 charge and are found on the outside of the nucleus Nucleus – made up of protons and neutrons, has an overall + charge

18. Atomic Numbers The atomic number of an element is the number of protons in the nucleus of an atom of that element. It is the number of protons that determines the identity of an element, as well as many of its chemical and physical properties. The number of protons for an element CANNOT be changed.

19. Atomic Numbers Because atoms have no overall electrical charge, the number of protons must equal the number of electrons. Therefore, the atomic number of an element also tells the number of electrons in a neutral atom of that element. The number of electrons can be changed when determining the charge of an ion.

20. Masses The mass of a neutron is almost the same as the mass of a proton. The sum of the protons and neutrons in the nucleus is the mass number of that particular atom. Isotopes of an element have different mass numbers because they have different numbers of neutrons, but they all have the same atomic number (number of protons & electrons)

21. ZXAZXA A = Mass # always a decimal # b/c it is an average of all isotopes Z = atomic # tells the Number of Protons # protons = # of electrons Neutrons = mass – atomic #

22. The # Of protons determines the identity of an element- this is represented by the letter Z If you change the # of electrons or neutrons without changing the number of neutrons the identity of the element remains the same

23. Isotopes – atoms of an element with the same number of protons but different number of neutrons. Ions – atoms of the same element with the same number of protons and a different number of electrons Ex: Na and its ion Na +

24. Isotopes When writing isotopes, the atomic number (or number of protons) will appear at the bottom left of the formula The mass number (number of protons plus neutrons) will appear at the top left of the formula. The element symbol will appear to the right of the numbers The different number of neutrons has NO bearing on chemical reactivity

25. Writing the Names of Isotopes When writing the name of an isotope, you will write the name of the element – the mass number For example 14 6 C would be named: Carbon - 14

26. Try the following NameSymbol# Protons# Neutrons# ElectronsMass # Carbon – 11 197 Au 79 12 2555 Oxygen - 15

27. Try the following NameSymbol# Protons# Neutrons# ElectronsMass # Carbon – 1111 C 6 65611 Gold - 197197 Au 79 11879197 Hydrogen – 3 3 H 1 1213 Manganese - 55 55 Mn 25 302555 Oxygen - 1515 O 8 87815

28. Atomic Mass Atomic mass –the weighted average mass of all the naturally occurring isotopes of that element. The number is usually located at the bottom of the periodic table and has decimal places

29. Calculating Atomic Mass

30. Copper exists as a mixture of two isotopes. The lighter isotope (Cu-63), with 29 protons and 34 neutrons, makes up 69.17% of copper atoms. The heavier isotope (Cu-65), with 29 protons and 36 neutrons, constitutes the remaining 30.83% of copper atoms.

31. Calculating Atomic Mass First, calculate the contribution of each isotope to the average atomic mass, being sure to convert each percent to a fractional abundance.

32. Calculating Atomic Mass The average atomic mass of the element is the sum of the mass contributions of each isotope.

33. Try this one… Calculate the atomic mass of germanium. 72.59 amu

34. You can tell many things from an isotope formula Hydrogen has three naturally occurring isotopes in nature: Hydrogen – 1, Hydrogen – 2, and Hydrogen – 3. –Which is the most abundant in nature? Hydrogen – 1 –Which is the heaviest? Hydrogen - 3

35. Quantum theory of the atom Bohr- said the lowest allowable energy state of an atom is called the ground state. Heisenberg’s uncertainty principle – it is impossible to know both the velocity and the momentum of an electron at the same time. Valence electrons- those found in the highest or outermost energy levels.

36. Bohr Bohr proposed that the hydrogen atom has only certain allowable energy states. Bohr suggested that the single electron in a hydrogen atom moves around the nucleus in only certain allowed circular orbits.

37. The Heisenberg Uncertainty Principle Heisenberg concluded that it is impossible to make any measurement on an object without disturbing the object—at least a little. The Heisenberg uncertainty principle states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.

38. Aufbau principle- says that electrons will always fill the lowest available energy levels first, i.e. 1s, 2s, etc. Pauli exclusion principle- a maximum of 2 electrons can occupy any given orbital at any time.

39. Aufbau Principle Each electron must occupy the lowest energy level first

40. Pauli Exclusion Principle In order for 2 electrons to share an orbital, they must have opposite spins In chemistry we designate spins with arrows. Therefore, if 2 electrons enter an orbital, they must enter 

41. Hund’s rule- that no 2 electrons can have the exact same 4 quantum #’s, that they will always have opposite spins. N, l, m, s

42. Hund’s Rule A single electron with the same spin must occupy each equal energy orbital before additional electrons will pair up with opposite spins You must fill before you pair For example, say you have 3 orbitals (don’t forget that each orbital can hold at most 2 electrons) You want to add 3 electrons

43. Orbital Notes Each element has energy levels represented with the symbol (n). Each element has sublevels determined by n 2 Each sublevel contains orbitals that can each contain a maximum of 2 e’.

44. Sublevels and Orbitals s sublevel = 1 orbital p sublevel = 3 orbitals d sublevel = 5 orbitals f sublevel = 7 orbitals

45. Quantum numbers – each electron has its own set of 4 quantum #’s. and no 2 e’ will have the same 4. They are represented with the symbols n, l, m, & s.

46. n = energy level = floor l = sublevel = apartment m= orbital = room s = spin = top or bottom

47. Possible Values for n, l, m n (shell) = 1, 2, 3, 4, … (whole numbers) L (sub shell) values from 0  (n -1) m L (orbital) values from – L to + L (including zero)

48. Quantum Numbers (n, l, m) L = Azimuthal Quantum Number Can have values from 0 to (n-1) for each value of n Defines the shape of the orbital L = 0  s L = 1  p L = 2  d L = 3  f Tells you what sub shell you are in

49. Quantum Numbers (n, l, m) m L = Magnetic Quantum Number Can have whole number values from – L to + L (including zero) This describes the orbital’s orientation in space (which axis it is on) Tells you what orbital you are in

50. Size

51. Shapes s p d f

52. Orientation

53. Examples What are the possible values for L if n =2? n = 2 Possible values for L go from 0 to (n-1) 0 to (2-1) 0 to 1 0, 1

54. Examples What are the possible values of n, L, and m in the 2s sub shell? n = 2 S means L = 0 m goes from – L to + L (-0 to + 0)  0 Therefore n = 2L = 0m = 0

55. Examples What are the possible values for n, L, & m in the 3d sub shell? n = 3 L = 2 M = -2  +2 (-2, -1, 0, 1, 2)

56. Example What is the designation for the sub shell where n = 2 and L = 1? 2p

57. Example What is the designation for the sub shell where n = 4 and L = 3? 4f

58. Summary Possible Values –L (0  n-1) –m (-L  +L) # of Possible Values –Orbitals (m) #m = 2L +1 #m = n 2 –Sub shells (L) # L = n –Electrons # electrons = 2n 2

59. Arrow Diagrams Lastly, you need to know the sequence of orbitals. I will NOT give you this on a test!

61. Arrow Diagrams Write the arrow diagram for sodium Arrows represent electrons, so before you can start, you need to find the number of electrons. How many electrons does sodium have? 11

62. Arrow Diagrams You will keep going until you reach 11 electrons Where do you always start? 1s 1s  Where to next? 2s 1s  2s 

63. Arrow Diagrams How many electrons have you used so far? 4 (only 7 more to go) What’s next? 2p 1s  2s  2p   

64. Arrow Diagrams Only one more left. Where does it go? 3s 1s  2s  2p    3s 

65. Arrow Diagrams Draw the arrow diagram for Br 1s  2s  2p    3s  3p    4s  3d      4p   ,

66. Electron Configurations Writing electron configurations is just a shorter way to write an arrow diagram You start with 1s and continue the configuration until you get the correct number of electrons

67. Electron Configurations Write the full electron configuration for K K (19) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 The number of electrons in each orbital are written as superscripts Remember s = 2, p = 6, d = 10, f = 14 Just follow the chart

68. Electron Configurations Write the full electron configuration for Kr Kr (36) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 Write the full electron configuration for P P (15) 1s 2 2s 2 2p 6 3s 2 3p 3

69. Noble Gas Configurations Noble Gas configuration is just a short hand way to write an electron configuration Steps 1.Find the element 2.Find the noble gas before that element (Group 8A) and place it in [brackets] 3.Move one spot 4.Start the configuration from there and keep going until you get to your element

70. Reading the periodic table s block – the first 2 columns of the periodic table (starts with 1s) p block – Groups 3A-8A, six columns (starts with 2p) d block – the center portion of the periodic table consisting of 10 columns (starts with 3d) f block – the two bottom rows of the periodic table consisting of 14 columns (starts with 4 f)

71. Noble Gas Configurations Write the noble gas configuration for Na Find the Noble Gas before Na [Ne] Move one spot and start the configuration from there [Ne]3s 1

72. Noble Gas Configurations Write the noble gas configuration for Br [Ar]4s 2 3d 10 4p 5 Write the noble gas configuration for Mn [Ar]4s 2 3d 5

73. Lewis Notes Write the symbol for the element involved. Determine the # of e’ in the outer energy level for the element from it’s electron configuration. Use dots to represent each outer e’ and place them in the appropriate sublevel using the following diagram:

74. p P X s » p