Topic 7. 1 Atomic Structure

1. Topic 7. 1 Atomic Structure

2. 7.1.1 Describe a model of the atom that features a small nucleus surrounded by electrons. • The modern atom has gone through a few stages of development • Dalton’s Atomic Therory – idea of an atom • JJ Thompson – 1890 – negative charge (electrons) • Earnest Rutherford – 1911 - positive nucleus (protons) • Niels Bohr – 1913 – orbital shells • Chadwick – 1932 – neutrons

3. 7.1.1 Describe a model of the atom that features a small nucleus surrounded by electrons. • This is a VERY simplified idea of the atom • Nucleus • Protons – positive charge – 1.6 x 10-19C • Neutrons – no charge • Diameter order of 10-15m • Electron “cloud” • Electrons – negative charge – 1.6 x 10-19C • Diameter order of 10-10m

4. 7.1.1 Describe a model of the atom that features a small nucleus surrounded by electrons. • The nucleus is about 100,000 times smaller than the electron orbits. • Imagine a pea in the center of a football field with the track being the orbits. • Protons and Neutrons have very similar mass. • Protons and Neutrons are about 1800 times bigger than electrons.

5. 7.1.2 Outline the evidence that supports a nuclear model of the atom Dalton’s Atomic Theory • All matter is composed of extremely small particles called atoms. • All atoms of a given element are identical. • Atoms cannot be created, divided into smaller particles, or destroyed. • Different atoms combine in simple whole number ratios to form compounds. • In a chemical reaction, atoms are separated, combined or rearranged.

6. Modern Atomic Theory Deomcritus Atoms Differences in atoms

7. Dalton’s Atomic Theory • All matter is composed of extremely small particles called atoms. • All atoms of a given element are identical. • Atoms cannot be created, divided into smaller particles, or destroyed. (This part proven wrong) • Different atoms combine in simple whole number ratios to form compounds. • In a chemical reaction, atoms are separated, combined or rearranged.

8. Modern Atomic Theory Deomcritus Atoms Differences in atoms • Dalton • Atoms • Sameness • Created/destroyed • Combination • Rearragement

9. J. J. Thomson – 1890-1900 Used cathode ray tube to prove existence of electron. Proposed “Plum Pudding Model” Cathode ray tube Stream of charged particles (electrons). http://www.youtube.com/watch?v=YG-Wz-arcaY http://www.youtube.com/watch?v=O9Goyscbazk Subatomic Particles and the Atom

10. Plum Pudding J. J. Thompson Plum Pudding Model Subatomic Particles and the Atom

11. Modern Atomic Theory Deomcritus Atoms Differences in atoms • Thompson • Atoms composed of electrons • Dalton • Atoms • Sameness • Created/destroyed • Combination • Rearragement

12. Gold Foil experiment Used to prove the existence of a positively charged core (Nucleus) Fired alpha particles(2protons and 2 neutrons) into very thin gold foil. The results were “like firing a large artillery shell at a sheet of paper and having the shell come back and hit you!” Ernest Rutherford

13. What should have happened Ernest Rutherford • What DID happened

14. After performing hundreds of tests and calculations, Rutherford was able to show that the diameter of the nucleus is about 105 times smaller than the diameter of the atom Ernest Rutherford

15. Modern Atomic Theory Deomcritus Atoms Differences in atoms • Thompson • Atoms composed of electrons • Rutherford • Positively Charged Nucleus • Dalton • Atoms • Sameness • Created/destroyed • Combination • Rearragement

16. Subatomic Particles and the Atom • Chadwick • Worked with Rutherford. • Noted there was energy in the nucleus, but wasn’t the protons. • Concluded that neutral particles must also exist in nucleus.

17. Subatomic Particles and the Atom • James Chadwick – 1932 • Bombarded a beryllium target with alpha particles • Alpha particles are helium nucleus • Discovered that , carbon was produced with another particle. • Concluded this particle had almost identical mass to proton but no charge. • Called it a neutron

18. Modern Atomic Theory Deomcritus Atoms Differences in atoms • Thompson • Atoms composed of electrons • Rutherford • Positively Charged Nucleus • Dalton • Atoms • Sameness • Created/destroyed • Combination • Rearragement • Chadwick • Neutrons exist in Nucleus

19. Subatomic Particles and the Atom • Three main particles: • Proton • Positive • In nucleus • Neutrons • Neutral • In nucleus • Electrons • Negative • Orbiting the nucleus (not inside)

20. 7.1.3 Outline one limitation of the simple model of the nuclear atom.7.1.4 Outline evidence for the existence of atomic energy levels. • If Rutherford’s was correct, electrons orbiting would undergo centripetal acceleration. • This would mean they would radiate electromagnetic waves. • Meaning they would loose energy • Meaning the atom would collapse on it’s self

21. 7.1.3 Outline one limitation of the simple model of the nuclear atom.7.1.4 Outline evidence for the existence of atomic energy levels. • If low-pressure gases are heated or current is passed through them they glow. • Different colors correspond to their wavelengths. • Visible spectrum 400nm(violet) to 750nm(red)

22. 7.1.3 Outline one limitation of the simple model of the nuclear atom.7.1.4 Outline evidence for the existence of atomic energy levels. • Gas – slit – slit – prism – viewing screen • When single element gases such as hydrogen and helium are excited only specific wave lengths were emitted. • These are called emission line spectra

23. 7.1.3 Outline one limitation of the simple model of the nuclear atom.7.1.4 Outline evidence for the existence of atomic energy levels. • Light – gas vapor – slit – slit – prism – viewing screen • If white light is pass through the gas the emerging light will show dark bands called absorption lines. • They correspond to the emission lines.

24. LIMITATION • Rutherford’s model didn’t explain why atoms emitted or absorbed only light at certain wavelengths. • 1885 JJ Balmer showed that hydrogen’s four emission lines fit a mathematical formula. • This “Balmer series” also show the pattern continued into non-visible ultra-violet and infra-red.

25. LIMITATION • Bohr called these “energy levels” • Reasoned that the electrons do not lose energy continuously but instead, lose energy in discrete amounts called “quanta”. • He agreed with Rutherford that electrons orbit the nucleus but only certain orbits were allowed.

26. LIMITATION • The electric force between protons and electrons holds electrons in orbit • Electron never found between these levels. (“jumps” instantly) • Only radiates energy when it “jumps” down. • Absorbs energy when it “jumps” up. • Total energy stays constant

27. LIMITATION • Bohr explained the emission and absorption line spectra with the idea that electrons absorbed only certain quantity of energy that allowed it to move to a higher orbit or energy level. • Each element has its own “finger print”.

28. Energy Level Diagram • Ground state – lowest energy level – smallest possible radius • Excited state – when an electron absorbs energy and jumps to a higher energy level. • Once an electron jumps back to a lower state it gives off energy in the form of a photon. • These photons are the emission spectrum.

29. Energy Level Diagram • The amount jumped correlates to the energy of the photon. • Greater the jump means the greater the energy is emitted. • Each jump corresponds to a different amount of energy being released. This means we can calculate the frequency and wavelength of light that will be produced.

30. Energy of a light quantum • E = hf • E = energy of a quantum • h = Planck’s constant (6.63 x 10-34Js) • f = frequency

31. Sample Problem C • An electron in a hydrogen atom drops from energy level E4 to energy level E2. What frequency of the emitted photon, and which line line in the emission spectrum corresponds to this event?

32. Sample Problem C • First find the amount of energy lost • Elost = E4 – E2 • Elost = (-0.85eV) – (-3.40eV) • Elost = 2.55 eV

33. Sample Problem C • Second, convert eV into J. • 1eV = 1.6 x 10-19J • Answer: 4.08 x 10-19J

34. Sample Problem C • Third use Planck’s equations to find the frequency. • E = hf • f = 6.15 x 1014 Hz

35. Sample Problem C • Fourth decide which line corresponds to this even. • Answer: Green light • v = f λ

36. Your practice • Practice C, pg 769 in book, #2 – 5

37. 7.1.5 Explain the terms nuclide, isotope and nucleon7.1.6 Define nucleon number A, proton number Z, and neutron number N. • Definitions • Nucleon – any of the constituents of a nucleus. Protons and neutrons. • Atomic Number – The number of protons in the nucleus. • Nucleon Number – The number of nucleons in the nucleus. AKA the mass number. (protons + neutrons) • Isotope – Nuclei which contain the same number of protons but different numbers of neutrons. • Nuclide – the nucleus of an atom. The nuclides of isotopes are different, even though they are the same element.

38. 7.1.5 Explain the terms nuclide, isotope and nucleon7.1.6 Define nucleon number A, proton number Z, and neutron number N. • Atomic Number (proton number), Z • How many protons there are. • This is what defines the element. • Ex. Hydrogen Z =1, Oxygen Z = 8 Carbon Z = 6 • Nucleon Number (mass number), A • How many nucleons there are. • Protons + neutrons • Number of neutrons, N • Mass number = atomic number + number of neutrons • A = Z + N

39. 7.1.5 Explain the terms nuclide, isotope and nucleon7.1.6 Define nucleon number A, proton number Z, and neutron number N. • Standard notation is: A over Z in front of element(X) *****Draw on board***** Isotopes • More evidence for neutrons is the existence of isotopes. • When nuclei of the same element have different numbers of neutrons. • Carbon has 6 isotopes: Carbon-11, Carbon-12, Carbon-13, Carbon-14, Carbon-15, Carbon-16. • All have 6 protons but each has different number of neutrons.

40. 7.1.5 Explain the terms nuclide, isotope and nucleon7.1.6 Define nucleon number A, proton number Z, and neutron number N. • The different isotopes don’t exist in nature in equal amounts. • Carbon: • C – 12 is most abundant (98.9%) • C – 13 is next (1.1%) • This is where atomic mass comes from. It’s the weighted average mass of all the different isotopes.

41. 7.1.5 Explain the terms nuclide, isotope and nucleon7.1.6 Define nucleon number A, proton number Z, and neutron number N. • Nuclei of different atoms are known as nuclides. • Ex. C – 12, C – 14 • Both are carbon but different isotopes • Their nuclei have different numbers of neutrons. • These are different nuclides.

42. 7.1.7 Describe the interactions in a nucleus • How do like charge (protons), stay stuck together? • We already know that like charges repel each other. • We have also seen that they are stronger than gravitational forces. • Strong Force – The force that binds the nucleus together. • It is an attractive force that acts between all nucleons. • Short– range interactions only (up to 10-15m)

46. 7.3.3 - Define the term unified atomic mass unit. • 7.3.4 - Apply the Einstein mass-energy equivalence relationship. • 7.3.5 - Define the concepts of mass defect, binding energy and binding energy per nucleon. • 7.3.6 - Draw and annotate a graph showing the variation with nucleon number of the binding energy per nucleon. • 7.3.7 - Solve problems involving mass defect and binding energy.

47. Unified Mass Unit • Because the mass of an atom is so small a new unit was created. • Some times called “Atomic mass unit” • 1 u = 1.66053886 x 10-27 kg • 12u = one atom of carbon-12

48. Resting Energy • Mass of a nucleus is sometimes expressed in terms of rest energy. • A particle has a certain amount of energy associated with its mass. Relationship between rest energy and mass: ER = mc2

49. Conservation of mass • It doesn’t always happen with nuclear processes. • Some times mass is converted or lost in the form of energy. • 1u = 931.49 MeV

50. Conservation of mass • So that means that one proton IS 938.3MeV of energy. • Mass is energy, energy is mass THEY ARE THE SAME THING!!! AHHHHHH!!!!!! Check out the table