Unit 2 – Atomic Structure & Nuclear Chemistry

1. Unit 2 – Atomic Structure & Nuclear Chemistry Part I – Atomic Theory, Subatomic Particles, and Average Atomic Mass

2. Part I Key Terms • Atomic mass - The mass of an atom of a chemical element expressed in atomic mass units. It is approximately equivalent to the number of protons and neutrons in the atom (the mass number) • Average atomic mass – Weighted average of all atoms of a particular element and is dependent on the mass of isotopes for an element and the relative population of each isotope • Bohr model - Devised by Niels Bohr, depicts the atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the nucleus • Dalton’s Postulates - States that matter is composed of extremely small particles called atoms; atoms are invisible and indestructable; atoms of a given element are identical in size, mass, and chemical properties; atoms of a specific element are different from those of another element; different atoms combine in simple whole-number ratios to form compounds; in a chemical reaction, atoms are separated, combined, or rearranged • Isotope -Atoms of the same element with different numbers of neutrons

3. Part I Key Terms (cont.) • Isotope notation - Subscripts and superscripts can be added to an element’s symbol to specify a particular isotope of the element and provide other important information. The atomic number is written as a subscript on the left of the element symbol, the mass number is written as a superscript on the left of the element symbol • Mass number - The total number of protons and neutrons in a nucleus. • Subatomic particles - The three kinds of particles that make up atoms: protons, neutrons, and electrons • Theory - An explanation supported by many experiments; is still subject to new experimental data, can be modified, and is considered valid if it can be used to make predictions that are proven true

4. Early Development of Atomic Theory • Major Contributors to Understanding Atomic Structure • Democritus – ancient Greek philosopher that originally stated all matter consists of atoms • 1605: Francis Bacon – published the scientific method • 1803: John Dalton – Postulates of Atomic Theory • 1897: J.J. Thomson – Discovery of the negatively charged electron and the mass to charge ratio of the electron • 1908 Robert Millikan – Determines the charge of the electron • 1911: Ernest Rutherford – Discovers positively charged nucleus • 1913: Niels Bohr – Theorizes structure of the electron cloud with energy levels and planetary orbits of electrons • 1932: James Chadwick – Discovers neutrons

5. Atomic Theory – John Dalton • John Dalton • English physicist • Experimented extensively with multiple gases and gaseous compounds • Contributions – Five Postulates of Atomic Theory • 1. All matter consists of tiny particles called atoms • 2. Atoms are indestructible and unchangeable. • 3. Elements are characterized by the mass of their atoms. • 4. When elements react, their atoms combine in simple, whole number ratios. • 5. When elements react, their atoms sometimes combine in more than one simple whole, number ratio.

6. Dalton’s Model of an Atom • He made no prediction about the construction of atoms believing them to be solid spheres. • Conclusions made based on his experiments and postulates: • Law of the Conservation of Mass – when chemical reactions occur, the atoms are only rearranged and there is no difference in mass following a chemical reaction • Law of Definite Proportions – elements combine in simple, low number ratios to form compounds (examples – H20, CO2) • Law of Multiple Proportions –elements combine in different simple, low number ratios to form different compounds (examples – H20 and H202; CO and CO2)

7. Atomic Theory – J.J. Thomson • Discovered the negatively charged electron and the mass to charge ratio of the electron • Used cathode ray tube • Beam of electrons deflected toward positive plate indicated the electron has negative charge • Amount of deflection indicates the mass to charge ratio

8. Thomson’s Experiment Image used courtesy of http://www.chemteam.info/AtomicStructure/Disc-of-Electron-Images.html

9. Thomson’s Model of the Atom Image used courtesy of http://www.kutl.kyushu-u.ac.jp/seminar/MicroWorld1_E/Part2_E/P24_E/Thomson_model_E.htm • Plum Pudding Model

10. Atomic Theory – Ernest Rutherford • Discovered positively charged nucleus • Used gold foil & detector ring • Fired alpha particles at foil which are positively charged • Most went through – atom mostly empty space • Some deflected – nucleus positively charged • Some bounced back – solid mass indicates nuclear core

11. Rutherford’s Experiment

12. Rutherford’s Model of the Atom Image used courtesy of http://www.bbc.co.uk/manchester/content/articles/2008/09/10/100908_rutherford_physics_feature.shtml • Nuclear Atomic Model

13. Atom Theory – Niels Bohr • Discovered electrons reside in energy levels with discrete amounts of energy • Mathematic modeling • Needed to explain why negatively charged electrons do not get absorbed into positively charged nucleus • Used information from Balmer, Lyman, & Paschen series • Emission spectra for Hydrogen explained by Rydberg equation

14. Bohr’s Model of the Atom Image used courtesy of http://www.blurtit.com/q982327.html • Electron Shell Model

15. 2 Regions of the Atom • Nucleus • Contains the protons and neutrons • Accounts for virtually all of the mass, but only a very small portion of the volume of the atom. • Has a positive charge equal to the number of protons. • Electron Cloud • Contains the electrons in orbitals • Has virtually no mass, but accounts for virtually all of the volume • Has a negative charge equal to the number of electrons.

16. Subatomic Particles • Electrons • Charge = -1 • Mass ≈ 0 amu • Location: in orbitals in the electron cloud (outside the nucleus) • Protons • Charge = +1 • Mass = 1 amu • Location: Inside the nucleus • Neutrons • Charge = 0 • Mass = 1 amu • Location: Inside the nucleus

17. Properties of the Atom • Mass • Measured in Atomic Mass Units (amu) • Equal to the sum of the number of protons and neutrons • Represented by the Mass Number • Charge • Neutral unless electrons gained or lost (ionized) • Number of electrons and protons is equal and, therefore balance out • Atomic Number • Equal to the number of protons • Define the element and its chemical properties

18. F 19 9 Symbology Example assuming neutral atom of Fluorine Atomic number: 9 Mass Number: 19 Protons: 9 Neutrons: 10 (mass number – atomic number) Electrons: 9

19. Atoms of the same element with different mass due to different number of neutrons Isotopes

20. Average Atomic Mass • Weighted average of all atoms of a particular element • Dependent on the mass of isotopes for an element and the relative population of each isotope • % mass oxygen-16: (15.99491) (.99759) = 15.9564 • % mass oxygen-17: (16.99913) (.00037) = 0.0063 • % mass oxygen-18: (17.99916) (.00204) = 0.0367 • Average Atomic Mass of Oxygen =15.9994

21. Naming Isotopes • Name of the element followed by the mass number of the isotope • Carbon – 12 = the name of the carbon atom with a mass number of 12 (6 protons and 6 neutrons) • Carbon – 14 = the name of the carbon atom with a mass number of 14 (6 protons and 8 neutrons) • Fluorine – 19 = the name of the Fluorine atom with a mass number of 19 (9 protons and 10 neutrons)

22. Energy Levels • Energy levels correspond to the energy of individual electrons. Each energy level has a discrete numerical value. • Different energy levels correspond to different numbers of electrons using the formula 2n2 where “n” is the energy level

23. Quantum Mechanical Model of Atomic Structure • 1900: Max Planck – Develops law correlating energy to frequency of light • 1905: Albert Einstein – Postulates dual nature of light as both energy and particles • 1924: Louis de Broglie – Applies dual nature of light to all matter • 1927: Werner Heisenberg – Develops Uncertainty Principle stating that it is impossible to observe both the location and momentum of an electron simultaneously • 1933: Erwin Schrodinger – Refines the use of the equation named after him to develop the concept of electron orbitals to replace the planetary motion of the electron

24. Orbitals • Impossible to determine the location of any single electron • Orbitals are the regions of space in which electrons can most probably be found • Four types of orbitals • s – spherically shaped • p – dumbbell shaped • d – cloverleaf shaped • f – shape has not been determined • Each additional energy level incorporates one additional orbital type • Each type of orbital can only hold a specific number of electrons

25. Orbital Types

26. Electron Configuration

27. Electron Configuration Notation • Find the element on the periodic table • Follow through each element block in order by stating the energy level, the orbital type, and the number of electrons per orbital type until you arrive at the element.

28. Samples of e- Configuration • Element Electron Configuration • H 1s1 • He 1s2 • Li 1s2 2s1 • C 1s2 2s2 2p2 • K 1s2 2s2 2p6 3s2 3p6 4s1 • V 1s2 2s2 2p6 3s2 3p6 4s2 3d3 • Br 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 (Note the overlap) • Pb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2

29. Noble Gas Electron Configuration Notation • Find element on the Periodic Table of Elements • Example: Pb for Lead • Move backward to the Noble Gas immediately preceding the element • Example: Xenon • Write symbol of the Nobel Gas in brackets • Example: [Xe] • Continue writing Electron Configuration Notation from the Noble Gas • Example: [Xe] 6s2 4f14 5d10 6p2

30. Valence Electrons • The electrons in the highest (outermost) s and p orbitals of an atom • The electrons available to be transferred or shared to create chemical bonds to form compounds • Often found in incompletely filled energy levels

31. Valence Electrons • Shortcut to finding valence electrons for main group elements • Family 1A (1) 1 valence electron • Family 2A (2) 2 valence electrons • Family 3A (13) 3 valence electrons • Family 4A (14) 4 valence electrons • Family 5A (15) 5 valence electrons • Family 6A (16) 6 valence electrons • Family 7A (17) 7 valence electrons • Family 8A (18) 8 valence electrons • Family3-12 have multiple possibilities and shortcuts do not work

32. Electron Dot Notation • Electron configuration notation using only the valence electrons of an atom. • The valence electrons are indicated by dots placed around the element’s symbol. • Used to represent up to eight valence electrons for an atom. One dot is placed on each side before a second dot is placed on any side. Valance Electrons: Sodium Magnesium Chlorine Neon 1 2 7 8 Electron Dot Notation: • • •• •• NaMg: Cl :: Ne : • • •• Oxidation Numbers: +1 +2 -1 0

33. Part II Key Terms • Alpha particle: A helium nucleus emitted by some radioactive substances • Beta particle: An energetic electron or positron produced as the result of a nuclear reaction or nuclear decay • Beta radiation: Radioactive decay in which an electron is emitted • Electron Configuration Notation -Consists of an element’s symbol, representing the atomic nucleus and inner-level electrons, that is surrounded by dots, representing the atom’s valence electrons. • Emission spectrum: The range of all possible wave frequencies of electromagnetic radiation, waves created by the systematic interactions of oscillating electric and magnetic fields • Energy Levels - A certain volume of space around the nucleus in which an electron is likely to be found. Energy levels start at level 1 and go to infinity. • Excited state: The state of an atom when one of its electrons is in a higher energy orbital than the ground state.

34. Part II Key Terms (cont.) • Gamma radiation: Electromagnetic radiation emitted during radioactive decay and having an extremely short wavelength • Ground state: The lowest energy state of an atom or other particle • Nuclear fission: Splitting of the nucleus into smaller nuclei • Nuclear fusion: Combining nuclei of light elements into a larger nucleus • Nucleon: a constituent (proton or neutron) of an atomic nucleus • Planck’s constant: As frequency increases, the energy of the wave increases • Radioactive decay: Spontaneous release of radiation to produce a more stable nucleus • Radioactive isotope: An isotope (an atomic form of a chemical element) that is unstable; the nucleus decays spontaneously, giving off detectable particles and energy

35. Electromagnetic (EM) Spectrum • The EM Spectrum is the range of all possible wave frequencies of electromagnetic radiation, waves created by the systematic interactions of oscillating electric and magnetic fields • The general term for all electromagnetic radiation is light • The range of the EM Spectrum is from very low frequency known as radio waves to very high frequency known as gamma radiation • The visible spectrum of light is in the center portion of this EM Spectrum • All EM Spectrum travels at the same speed in a vacuum – this speed is known as the speed of light, 3.00 x 108 m/s

36. EM Spectrum Image used courtesy of http://9-4fordham.wikispaces.com/Electro+Magnetic+Spectrum+and+light

37. Speed of Light and Frequency • Since the speed of all EM radiation is the same, there is a clear mathematical relationship between the frequency of the light and its wavelength • All waves travel at a speed that is equal to the product of its frequency (the reciprocal of time) and its wavelength (distance) c = f λ • The speed of EM radiation is fixed at 3.00 x 108 m/s • Therefore: 3.00 x 108 m/s = f λ Speed of light = frequency x wavelength • As frequency increases, wavelength decreases. As wavelength increases, frequency decreases • Example: If frequency doubles, wavelength is cut in half

38. As f ↑, λ↓: Calculations • If the wavelength of a radio wave is 15 meter, what is its frequency? 3.00 x 108 m/s = f (10 m) (3.00 x 108 m/s) / 15 m = f 2.0 x107 s-1 = f Frequency = 2.0 x107 Hertz • If the frequency of gamma radiation is 6.25 x 1022 Hertz, what is its wavelength? 3.00 x 108 m/s = (6.25 x 1022 s-1) λ (3.00 x 108 m/s) / (6.25 x 1022 s-1) = λ 4.80 x10-15 m= f Wavelength = 4.80 x10-15 m

39. Planck’s Law • Max Planck determined in 1900 there was a mathematical relationship between the energy of EM radiation and the frequency of that radiation: As frequency increases, the energy of the wave increases E = h f Energy = Planck’s constant x frequency E = (6.63 x 10-34 Joule seconds) f

40. Planck’s Law Calculations • Example: If the wavelength of green light is 5.21 x 10-7 meters, what is the energy of this light? 3.00 x 108 m/s = f (5.21 x 10-7 m) (3.00 x 108 m/s) / 5.21 x 10-7 m = f 5.76 x1014 s-1 = f Frequency = 5.76 x1014 Hertz E = (6.63 x 10-34 Joule seconds) (5.76 x1014 s-1) E = 3.82 x10-19 Joules

41. Implication of Planck’s Law • In order to move an electron to a higher energy level, excite an electron, energy must be absorbed to move the electron • Since electrons exist in fixed energy levels with a specific amount of energy, the amount of energy needed is a finite amount equal to the difference in the energy associated with the ground state of the electron and the energy associated with the level to which the electron is excited • If the energy related to the excited electron is removed, the electron will return to its ground state and the energy released is equal to the energy absorbed to excite it • The energy released is released as light • The overall result is that every element has a unique spectra of light associated with it and the spectra can be used to identify the element

42. Nuclear Reactions • All nuclear reactions are based on Einstein’s Theory of Relativity • At speeds approaching the speed of light, energy and mass are interchangeable • E = mc2 Energy = mass x (speed of light)2 • Mass can be converted to energy and vice versa

43. Mass Defect • There is a difference between the mass of an atom and the various particles that make up the atom • This difference is called the mass defect of the atom • This mass defect is the binding energy of the atom • In nuclear reactions, the binding energy is released as energy (heat, light, or gamma radiation) and/or particles with measureable mass

44. Types of Nuclear Reactions • Fission – Splitting of the nucleus into smaller nuclei • Fusion – Combining nuclei of light elements into a larger nucleus • Radioactive Decay – Spontaneous release of radiation to produce a more stable nucleus

45. Fission • Nucleus splits into smaller nuclei when struck by a neutron of sufficient energy • Tremendous release of energy • When controlled can produce huge amounts of power in nuclear reactors • Naturally occurs in uranium and other ores in spontaneous fission • Clean source of energy with no carbon footprint • Produces radioactive nuclear waste with long term environmental and health considerations

46. Fission Process

47. Fission and Nuclear Reactors

48. Fusion • Lighter nuclei (such as hydrogen) combined to form heavier nuclei • Tremendous release of energy • 2H + 3H 4He + 1n + energy Deuterium Tritium Helium (occurs naturally in water) • Powers the sun and stars • No practical application to produce usable energy at this time

49. Fusion Process

50. Radioactive Decay • Spontaneous release of radiation by unstable nuclei in order to increase stability • Radiation can be either energy alone (gamma) or energy accompanied by release of a particle (all of the other forms of decay)