1. Unit 4: Atomic Theory Structure of the Atom (& Radioactivity)

2. Early Atomic Theories

3.  Atomists and Democritus Greeks approx. 2,500 years ago Matter was made up of atoms  “atomos” or “Indivisible” particles Seashell experiment—broken into smaller & smaller pieces

4. Early Atomic Theories  John Dalton 1766-1844; returned to theory of atoms Atoms are like billiard balls (solid spheres) which can’t be broken down further 4 major postulates 1. All elements are composed of atoms 2. Atoms of the same element are identical 3. Atoms can physically mix or chemically combine in simple whole number ratios 4. Reactions occur when atoms separate, join, or rearrange

5. Dalton’s Model of the Atom No subatomic particles!

6. Early Atomic Theories  William Crookes Developed Crookes tube (a.k.a cathode ray tube) in 1870’s First evidence for existence of electrons because you could “see” electrons flow and confirm their existence. Tube is precursor to today’s TV picture tubes

7. Building the Atom – The Electron  J.J. Thomson Discovered electron in 1897 Discovered positively charged particles surrounded by electrons Found the ratio of the charge of an electron to its mass to be 1/1837

8. Thomson’s Cathode Ray Tube Experiment  J.J. Thomson Video: Cathode Ray Tube Demo

9. Building the Atom – The Electron  J.J. Thomson Cathode ray tube experiments – advancement of Crookes tube “plum-pudding model”

10. Thomson’s Model of the Atom (Plum Pudding Model)

11. Millikan’s Oil Drop Experiment  Robert Millikan Oil drop experiment Determined the charge and mass of an electron Video: Millikan's Oil Drop Experiment

12. Building the Atom – The Nucleus  Ernest Rutherford Discovered nucleus (dense core of atom) in 1911 Famous Gold foil experiment Quote from E.R.’s Lab Notebook  “It is about as incredible as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you.” -ER

13. Rutherford’s Gold Foil Experiment Video Clip: Rutherford Gold Foil Experiment

14. Rutherford’s Model of the Atom

15. Building the Atom – The Neutron  James Chadwick Discovered the neutron (no charge, but same mass as proton) Neutrons help disperse the strong repulsion of positive charges  Nucleus diameter = 10 -5 nm  Atom diameter = 10 -1 nm  If Nucleus = basketball --> then Atom = 6 miles wide!

16. Building the Atom  Niels Bohr Improved on Rutherford’s work “Planetary model”- positive center is surrounded by electrons in defined orbits circling the center

17. Bohr Model of the Atom (Planetary Model)

18. Bohr Model of the Atom Vocab.  Energy level – the location where an electron is found at a set distance from the nucleus dependent on the amount of energy it has  Ground state – the typical energy level where an electron is found; lowest energy  Excited state – an energy level higher than the ground state for an electron; temporary condition

19. Schrödinger Model (Quantum Mechanical Model)  Quantum Mechanical Model Erwin Schrödinger; Mathematical model Electron locations are based on probability Electrons are not particles, but waves!  Interactive Simulation: try it! Interactive Simulation: try it! Defined: Orbital – region where an electron is likely to be found 90% of the time

20. Schrödinger Model of the Atom (Quantum Mechanical Model)

21. Atomic Theory  Atom – the smallest particle of matter that retains its properties. Smallest individual unit of an element  One atom of hydrogen is different from one atom of carbon.  Subatomic particles – the component parts of an atom: proton, neutron, and electron

22. Atomic Theory  Ion - atom with the same number of protons but a different number of electrons i.e. an atom with a charge! If the atom has a (+) charge it has more protons than electrons. If the atom has a (-) charge it has more electrons than protons.

23. Subatomic Particles Subatomic Particle Mass and Abbreviation ChargeLocationDiscoverer Proton p + Mass =1 amu +1Nucleus---- Neutron n Mass =1 amu 0Nucleus Chadwick in 1932 Electron e - Mass ≈ zero amu Electron cloud (outside nucleus) Thomson in 1897

24. Atomic Symbols Atomic symbol – the letter or letters that represent an element. 13 Al 26.981 Atomic number Atom symbol Atomic mass or weight

25. Atomic Number Atomic number = the number of protons in the nucleus. (same for every atom of that element) 13 Al 26.981 Atomic number Atom symbol Atomic mass or weight

26. Mass Number Mass Number = # protons + # neutrons  A Boron atom can have: 5 p + 5 n = 10 amu Named as boron-10 Mass number Atomic number

27. Calculations w/ Subatomic Particles  Atomic number = # of protons  Mass number = # of protons + # of neutrons  (For a neutral atom): # of protons = # of electrons  (For a charged ion): Charge = #p + - #e -

28. Isotope Notation  Isotope (Isotopic Notation) Mass # Atomic # Atomic Symbol Z A X Example: Uranium-238

29. Example 

30. 

31. Sample Problem Write the atomic symbols for the following:  The isotope of carbon with a mass of 13  The nuclear symbol when A = 92 and the number of neutrons = 146. 31

32. Isotopes  Isotope – atoms of the same element with different numbers of neutrons (different mass numbers) Example: Carbon-12 Carbon-14  Atomic mass – weighted average of the masses of all the isotopes of an element

33. Atomic Mass The weighted average is the addition of the contributions from each isotope. Isotopic Abundance is the percent or fraction of each isotope found in nature. 33

34. Most Abundant Isotope 34 13 Al 26.981 Atomic number Atom symbol Atomic mass or weight Usually can round atomic mass on the periodic table to nearest whole number (but not always!!)

35. Example: Determine the average atomic mass of magnesium which has three isotopes with the following masses: 23.98 amu (78.6%), 24.98 amu (10.1%), 25.98 amu (11.3%). 1) Multiply the mass number of the isotope by the decimal value of the percent for that isotope. 2) Add the relative masses of all of the isotopes to get the atomic mass of the element. 35

36. Example: Determine the average atomic mass of magnesium which has three isotopes with the following masses: 23.98 amu (78.6%), 24.98 amu (10.1%), 25.98 amu (11.3%). 36

37. Now You Try! If 90% of the beryllium in the world has a mass number of 9 and only 10% has a mass number of 10, what is the atomic mass of beryllium?

38. Radioactivity - Vocabulary  Radioactivity - the spontaneous emission of radiation from a substance  Radiation - rays and/or particles emitted from radioactive material  Nuclear reactions - reactions involving changes in an atom’s nucleus

39. Radioactivity  Radioactive isotopes are unstable These isotopes decay over time by emitting particles and are transformed into other elements  Particles emitted: Alpha (α) particles: helium nuclei Beta (β) particles: High speed electrons Gamma (γ) rays: high energy light 39

40. Types of Radiation – α particles  Alpha radiation - stream of high energy alpha particles Alpha particles consist of 2 protons and 2 neutrons and are identical to helium-4 nucleus. Symbol: 4 He 2+ 2 Not much penetrating power, travel a few centimeters, stopped by paper, no health hazard

41. Types of Radiation – β particles  Beta radiation - High speed electrons To form beta radiation a neutron splits into a proton and an electron The proton stays in nucleus and the electron propels out at high speed.  Symbol: 0 e- 0 β -1 -1  100 times more penetrating then alpha radiation, pass through clothing to damage skin

42. Types of Radiation –  Radiation  Gamma radiation Similar to X rays Doesn’t consist of particles (instead, high energy light)  Symbol: 0  0 Penetrates deeply into solid material, body tissue, stopped by Pb or concrete, dangerous

43. End of Unit 4!