2. Atomic Structure Chemistry

3. The Greeks History of the Atom Not the history of atom, but the idea of the atom In 400 B.C the Greeks tried to understand matter (chemicals) and broke them down into earth, wind, fire, and air. Democritus and Leucippus Greek philosophers  

4. The Hellenic Market FireWater Earth Air ~ ~

5. ~400 BCDemocritus First suggested the existence of atoms. Name atom - atomos All matter is made up of small indestructible particles called atoms

6. DEMOCRITUS (400 BC) – First Atomic Hypothesis Atomos: Greek for “uncuttable”. Chop up a piece of matter until you reach the atomos. Properties of atoms: indestructible. changeable, however, into different forms. an infinite number of kinds so there are an infinite number of elements. hard substances have rough, prickly atoms that stick together. liquids have round, smooth atoms that slide over one another. smell is caused by atoms interacting with the nose – rough atoms hurt. sleep is caused by atoms escaping the brain. death – too many escaped or didn’t return. the heart is the center of anger. the brain is the center of thought. the liver is the seat of desire. “Nothing exists but atoms and space, all else is opinion”. Democritus

7. Aristotle and the Philosophers The opposing Philosopher to Democritus was Aristotle. He believed that matter could be continuously broken in half, and that there was no end to this series. However, both men were philosophers and had no concrete information to back up their opinion

8. Scientist Enter the Picture (1627-1691) Robert Boyle (England) extended mathematics to chemistry and revived atomic theory. 1661 – published “The Skeptical Chymist”

9. AntoineJoseph LavoisierProust (1789) Antoine Lavoisier (France) demonstrated the conservation of matter (matter can be neither created nor destroyed) in a chemical reaction and defined the difference between an element and a compound. (1800)Law of Definite Proportions – The proportion by mass of elements in a compound is always the same.

10. Daltons Atomic Theory 1766-1844 Dalton stated that elements consisted of tiny particles called atoms He also called the elements pure substances because all atoms of an element were identical and that in particular they had the same mass.

11. Dalton’s Atomic Theory 1803 First atomic theory based on proof. 1. All matter is made up of atoms. 2. Atoms of the same element are the same, atoms of different elements are different. 3. Atoms combine in small whole number ratios to form compounds. 4. Atoms are rearranged in a chemical reaction.

12. Dalton’s Symbols John Dalton 1808

13. Daltons’ Models of Atoms Carbon dioxide, CO 2 Water, H 2 O Methane, CH 4

14. Date: 1806; Author: John Dalton; Published in: New System of Chemical PhilosophyJohn Dalton

15. Structure of Atoms Scientist began to wonder what an atom was like. Was it solid throughout with no internal structure or was it made up of smaller, subatomic particles? It was not until the late 1800’s that evidence became available that atoms were composed of smaller parts.

16. History: On The Human Side 1834 Michael Faraday - electrolysis experiments suggested electrical nature of matter 1895 Wilhelm Roentgen - discovered X-rays when cathode rays strike anode 1896 Henri Becquerel - discovered "uranic rays" and radioactivity 1896 Marie (Marya Sklodowska) and Pierre Curie - discovered that radiation is a property of the atom, and not due to chemical reaction. (Marie named this property radioactivity.) 1897 Joseph J. Thomson - discovered the electron through Crookes tube experiments 1898 Marie and Piere Curie - discovered the radioactive elements polonium and radium 1899 Ernest Rutherford - discovered alpha and beta particles 1900 Paul Villard - discovered gamma rays 1903 Ernest Rutherford and Frederick Soddy - established laws of radioactive decay and transformation 1910 Frederick Soddy - proposed the isotope concept to explain the existence of more than one atomic weight of radioelements 1911 Ernest Rutherford - used alpha particles to explore gold foil; discovered the nucleus and the proton; proposed the nuclear theory of the atom 1919 Ernest Rutherford - announced the first artificial transmutation of atoms 1932 James Chadwick - discovered the neutron by alpha particle bombardment of Beryllium 1934 Frederick Joliet and Irene Joliet Curie - produced the first artificial radioisotope 1938 Otto Hahn, Fritz Strassmann, Lise Meitner, and Otto Frisch - discovered nuclear fission of uranium-235 by neutron bombardment 1940 Edwin M McMillan and Philip Abelson - discovered the first transuranium element, neptunium, by neutron irradiation of uranium in a cyclotron 1941 Glenn T. Seaborg, Edwin M. McMillan, Joseph W. Kennedy and Arthur C. Wahl - announced discovery of plutonium from beta particle emission of neptunium 1942 Enrico Fermi - produced the first nuclear fission chain-reaction 1944 Glenn T. Seaborg - proposed a new format for the periodic table to show that a new actinide series of 14 elements would fall below and be analogous to the 14 lanthanide-series elements. 1964 Murray Gell-Mann hypothesized that quarks are the fundamental particles that make up all known subatomic particles except leptons.

17. Subatomic Particles  Today… Dalton’s Atomic Theory has one important change  Atoms are divisible…  3 types of subatomic particles ProtonsProtons NeutronsNeutrons ElectronsElectrons

18. Electrons  J.J. Thomson (English 1897) – Identified the 1 st subatomic particleIdentified the 1 st subatomic particle Cathode-ray tubeCathode-ray tube

19. J.J. Thomson He proved that atoms of any element can be made to emit tiny negative particles. From this he concluded that ALL atoms must contain these negative particles. He knew that atoms did not have a net negative charge and so there must be balancing the negative charge. J.J. Thomson

20. Thomson’s Electron Model  Named electrons  Symbol: e -  Charge: (-1)  Mass: ~ 0 amu  Plum Pudding/ Chocolate Chip Cookie

21. Plum Pudding Model.

22. William Thomson (Lord Kelvin) In 1910 proposed the Plum Pudding model –Negative electrons were embedded into a positively charged spherical cloud. Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 56 Spherical cloud of Positive charge Electrons

23. Thomson Model of the Atom J.J. Thomson discovered the electron and knew that electrons could be emitted from matter (1897). William Thomson proposed that atoms consist of small, negative electrons embedded in a massive, positive sphere. The electrons were like currants in a plum pudding. This is called the ‘plum pudding’ model of the atom. - electrons - - - - - - -

24. Millikans oil drop experiment (1909) Determined the charge of an electron.

25. Lord Earnest Rutherford 1911 Gold Foil Experiment Nuclear model of the atom

26. Ernest Rutherford (1871-1937) Learned physics in J.J. Thomson’ lab. Noticed that ‘alpha’ particles were sometime deflected by something in the air. Gold-foil experiment Rutherford PAPER Rutherford PAPER Animation by Raymond Chang Animation by Raymond Chang – All rights reserved.

27. Niels Bohr (1913) Planetary model of the atom Electrons have certain energies which allow them to stay in certain orbits around the nucleus.

28. Erwin Schrodinger 1926 Developed a mathematical equation to describe the motion of electrons in atoms. Quantum Mechanical Model of the atom

29. Electron Cloud Model

30. James Chadwick (1932) In 1930 it was discovered that Beryllium, when bombarded by alpha particles, emitted a very energetic stream of radiation. Like gamma rays, these rays were extremely penetrating and since they were not deflected upon passing through a magnetic field, neutral. In 1932, James Chadwick proposed that this particle was a neutron.

31. Summary Democritus proposed the existence of atoms around 400 BC. John Dalton introduced a new form of the ancient Greek idea of atoms at the beginning of the nineteenth century. In 1897, J.J. Thomson discovered the electron and suggested the 'plum pudding' model of the atom. In 1911, Rutherford suggested that the atom had a dense positive nucleus surrounded by the electrons in empty space In 1914, Bohr modified Rutherford's model by introducing the idea of energy levels. The electrons orbit the nucleus like planets around the sun. In 1926 Schrodinger developed the mathematical model now called the electron cloud model.

32. Atomic Structure

33. Putting the Pieces Together Where is all of the atom’s mass located? What overall charge does the nucleus have?

34. Putting the Pieces Together What charge do the energy levels have? What is required for an atom to be neutrally charged?

35. Atomic Number – Number of Protons Element Symbol – Is either one capital letter, or one capital and one lower case letter. Average Atomic Mass – Element Name

36. Mass Number If… Mass Number = # protons + # neutrons …how could you calculate the number of neutrons given the mass number and # protons? # neutrons = Mass Number - # protons

37. Atomic Structure - Notations Hyphen Notation Given: Element name use to find p + Mass number use to find n 0 Nuclear Symbol Given: Mass number (superscript) Atomic number (subscript) Element symbol Carbon - 13

38. Recap… Isotope Nuclear Symbol Atomic # Mass # # Protons # Electrons # Neutrons Carbon-12 Boron - 11 Xenon - 131

39. Drawing the Bohr Model of the Atom

40. Models of the Atom Dalton’s model (1803) Thomson’s plum-pudding model (1897) Rutherford’s model (1909) Bohr’s model (1913) Charge-cloud model (present) Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 125 Greek model (400 B.C.) 1800 1805..................... 1895 1900 1905 1910 1915 1920 1925 1930 1935 1940 1945 1803 John Dalton pictures atoms as tiny, indestructible particles, with no internal structure. 1897 J.J. Thomson, a British scientist, discovers the electron, leading to his "plum-pudding" model. He pictures electrons embedded in a sphere of positive electric charge. 1904 Hantaro Nagaoka, a Japanese physicist, suggests that an atom has a central nucleus. Electrons move in orbits like the rings around Saturn. 1911 New Zealander Ernest Rutherford states that an atom has a dense, positively charged nucleus. Electrons move randomly in the space around the nucleus. 1913 In Niels Bohr's model, the electrons move in spherical orbits at fixed distances from the nucleus. 1924 Frenchman Louis de Broglie proposes that moving particles like electrons have some properties of waves. Within a few years evidence is collected to support his idea. 1926 Erwin Schrödinger develops mathematical equations to describe the motion of electrons in atoms. His work leads to the electron cloud model. 1932 James Chadwick, a British physicist, confirms the existence of neutrons, which have no charge. Atomic nuclei contain neutrons and positively charged protons. + - - - - - e e e + + + + + + + + e ee e e e e

41. Models of the Atom Dalton’s model (1803) Thomson’s plum-pudding model (1897) Rutherford’s model (1909) Bohr’s model (1913) Charge-cloud model (present) Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 125 Greek model (400 B.C.)

42. II. The electron as a wave Schrödinger’s wave equation – Used to determine the probability of finding the H electron at any given distance from the nucleus – Electron best described as a cloud Effectively covers all points at the same time (fan blades)

43. Quantum Mechanics Schrödinger Wave Equation Schrödinger Wave Equation (1926) quantized – finite # of solutions  quantized energy levels probability – defines probability of finding an electron Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Erwin Schrödinger ~1926

44. Quantum Mechanics Orbital (“electron cloud”) – Region in space where there is 90% probability of finding an electron Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Electron Probability vs. Distance Electron Probability (%) Distance from the Nucleus (pm) 100150200250500 0 10 2020 30 40 Orbital 90% probability of finding the electron

45. Quantum Numbers UPPER LEVEL Four Quantum Numbers: – Specify the “address” of each electron in an atom Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

46. III. Quantum Numbers Used the wave equation to represent different energy states of the electrons Set of four #’s to represent the location of the outermost electron Here we go…

47. Quantum Numbers Principal Quantum Number n Principal Quantum Number ( n ) Angular Momentum Quantum # l Angular Momentum Quantum # ( l ) Magnetic Quantum Number m l Magnetic Quantum Number ( m l ) Spin Quantum Number Spin Quantum Number ( m s )

48. Quantum Numbers Principal Quantum Number n 1. Principal Quantum Number ( n ) – Energy level – Size of the orbital – n 2 = # of orbitals in the energy level Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem 1s1s 2s2s 3s3s

49. Quantum Numbers s p d f Angular Momentum Quantum # l 2. Angular Momentum Quantum # ( l ) – Energy sublevel – Shape of the orbital Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

50. Shapes of s, p, and d-Orbitals s orbital p orbitals d orbitals

51. s, p, and d-orbitals A s orbitals: Hold 2 electrons (outer orbitals of Groups 1 and 2) B p orbitals: Each of 3 pairs of lobes holds 2 electrons = 6 electrons (outer orbitals of Groups 13 to 18) C d orbitals: Each of 5 sets of lobes holds 2 electrons = 10 electrons (found in elements with atomic no. of 21 and higher) Kelter, Carr, Scott,, Chemistry: A World of Choices  1999, page 82

52. Copyright © 2006 Pearson Benjamin Cummings. All rights reserved.

53. Maximum Capacities of Subshells and Principal Shells n 1 2 3 4...n l 0 0 1 0 1 2 0 1 2 3 Subshell designation designation s s p s p d s p d f Orbitals in subshell subshell 1 1 3 1 3 5 1 3 5 7 Subshell capacity capacity 2 2 6 2 6 10 2 6 10 14 Principal shell capacity capacity 2 8 18 32...2n 2 Hill, Petrucci, General Chemistry An Integrated Approach  1999, page 320

54. Quantum Numbers Magnetic Quantum Number m l 3. Magnetic Quantum Number ( m l ) – Orientation of orbital – Specifies the exact orbital within each sublevel Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

55. Quantum Numbers Spin Quantum Number 4. Spin Quantum Number ( m s ) – Electron spin  +½ or -½ – An orbital can hold 2 electrons that spin in opposite directions. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

56. Copyright © 2006 Pearson Benjamin Cummings. All rights reserved.

57. Quantum Numbers 1. Principal #  2. Ang. Mom. #  3. Magnetic #  4. Spin #  energy level sublevel (s,p,d,f) orbital electron Pauli Exclusion Principle Pauli Exclusion Principle – No two electrons in an atom can have the same 4 quantum numbers. – Each electron has a unique “address”: Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Wolfgang Pauli

58. Level n 1 2 3 Sublevel l Orbital m l Spin m s 00 00 1 0 0 1 0 21 0 -2 2101 = +1/2 = -1/2 Allowed Sets of Quantum Numbers for Electrons in Atoms

59. Feeling overwhelmed? Read Section 5.10 - 5.11! Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem "Teacher, may I be excused? My brain is full." Chemistry

60. Electron Orbitals: Electron orbitals Equivalent Electron shells (a) 1s orbital (b) 2s and 2p orbitalsc) Neon Ne-10: 1s, 2s and 2p 1999, Addison, Wesley, Longman, Inc.

61. Electron Configuration

62. H = 1s 1 1s1s He = 1s 2 1s1s Li = 1s 2 2s 1 1s1s 2s2s Be = 1s 2 2s 2 1s1s 2s2s C = 1s 2 2s 2 2p 2 1s1s 2s2s 2px2px 2py2py 2pz2pz S = 1s 2 2s 2 2p 6 3s 2 3p 4 1s1s 2s2s 2px2px 2py2py 2pz2pz 3s3s 3px3px 3py3py 3pz3pz THIS SLIDE IS ANIMATED IN FILLING ORDER 2.PPTFILLING ORDER 2.PPT

63. H = 1s 1 1s1s He = 1s 2 1s1s Be = 1s 2 2s 2 1s1s 2s2s +1 e-e- +2 e-e- e-e- +4 e-e- e-e- e-e- e-e- Coulombic attraction holds valence electrons to atom. Valence electrons are shielded by the kernel electrons. Therefore the valence electrons are not held as tightly in Be than in He.

64. Orbital Filling Element 1s 2s 2p x 2p y 2p z 3s Configuration Orbital Filling Element 1s 2s 2p x 2p y 2p z 3s Configuration Electron Configurations Electron H He Li C N O F Ne Na 1s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 1s 2 2s 2 2p 5 1s 2 2s 2 2p 4 1s 2 2s 2 2p 3 1s 2 2s 2 2p 2 1s 2 2s 1 1s 2 NOT CORRECT Violates Hund’s Rule Electron Configurations Electron H He Li C N O F Ne Na 1s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 1s 2 2s 2 2p 5 1s 2 2s 2 2p 4 1s 2 2s 2 2p 3 1s 2 2s 2 2p 2 1s 2 2s 1 1s 2

65. Filling Rules for Electron Orbitals Aufbau Principle: Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for. Pauli Exclusion Principle: An orbital can hold a maximum of two electrons. To occupy the same orbital, two electrons must spin in opposite directions. Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results. *Aufbau is German for “building up”

66. Order in which subshells are filled with electrons 1s2s3s4s5s6s7s1s2s3s4s5s6s7s 2p3p4p5p6p 2p3p4p5p6p 3d4d5d6d 3d4d5d6d 4f5f 4f5f 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d … 2 2 6 2 6 2 10 6 2 10

67. 4f4f 4d4d 4p4p 4s4s n = 4 3d3d 3p3p 3s3s n = 3 2p2p 2s2s n = 2 1s1s n = 1 Energy Sublevels s s s s p p p d df 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 …

68. RIGHT WRONG General Rules Hund’s Rule Hund’s Rule – Within a sublevel, place one electron per orbital before pairing them. – “Empty Bus Seat Rule” Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

69. O 8e - Orbital Diagram Electron Configuration 1s 2 1s 2 2s 2 2s 2 2p 4 Notation 1s 2s 2p Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem O 15.9994 8

70. Shorthand Configuration S 16e - Valence Electrons Core Electrons S16e - [Ne] 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 4 Notation Longhand Configuration Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem S 32.066 16