Chapter 4

1. Chapter 4 Atoms

2. Section 1: The Development of Atomic Theory pp. 113 - 118

3. Beginnings of Atomic Theory • First theory that was proposed2,000years ago b. Democritus suggested that the universe was made of indivisible units he called atoms. c. He did not have evidence to convince people of his theory

4. B. Dalton’s Atomic Theory a. Proposed in 1808byJohn Dalton (English schoolteacher) b. All atoms of a given element were exactly alike, and atoms of differentelementscouldjointo form compounds

5. C. Thomson’s Model of the Atom a. Developed in 1897 by J.J. Thomson (British scientist) b. Cathode rays: mysterious rays in vacuum tubes c. Thomson was studying these cathode rays and it was discovered that atoms were not indivisible d. Cathode ray tube experiment suggested that cathode rays were made of negativelycharged particles that came frominside atoms i. Atoms could be divided into smaller parts!

6. Cathode Ray Experiment

7. D. Rutherford’s Model of the Atom a. Ernest Rutherford (British scientist) b. Proposed that most of the massof the atom wasconcentratedat the atom’s center c. Atom’s positive charge is concentrated in the centerof the atom(nucleus) andnegative electrons orbit

8. Rutherford's Experiment

10. Section 2: The Structure of the Atom pp. 119 - 127

11. What is in an atom?? a. Three main subatomic particles are distinguished by: - mass - charge - locationin the atom b. Nucleus: center of each atom, small and dense, made of protons and neutrons

12. c. Protons: positive charge d. Neutrons: no charge e. Electrons: move around outside the nucleus in a cloud f. Mass of electron is muchsmallerthan that of a protonorneutron

15. g. Each element has auniquenumber of protons i. elements are defined by the number of protonsin an atomof that element h. Most atoms have an equalnumber of protonsand electrons, whose chargescancel out

16. Example: He (helium) atom Charge of two protons: +2 Charge of two neutrons: 0 Charge of two electrons: -2_ 0

17. ii. If an atom gains or loses electrons, it becomescharged (ion)

18. Electric force holdsatom together i. Electric force: positive and negative charges that attract each other

19. B. Atomic Number and Mass Number a. Atoms of each elementhave the same number of protons, but they can have differentnumbers of neutrons

20. b. Atomic Number (Z): number of protons i. Each element has a unique atomic number ii, hydrogen has only one proton, so Z=1 iii. Uranium has 92 protons, Z = 92 iv. Atomic number of an atom NEVERchanges

21. c. Mass number (A): number of protons plus the number of neutrons i. Fluorine has 9 protons and 10 neutrons, so A=19 ii. Oxygen has 8protons and 8 neutrons, A=16

22. iv. Although atoms of an element have the same atomic number, they can have differentmass numbers because the numberof neutrons canvary.

23. What is Lithium’s A? What is Lithium’s Z?

24. C. Isotopes a. Isotope: atom that has the same number of protons but a different number of neutrons relative to other atoms of the same element b. have different masses because numbers of neutrons differs

26. Protium c. Hydrogen isotopes: i. Protium: A=1, Z=1 ii.Deuterium: A=2, Z =1 iii. Tritium: A=3, Z=1

27. Protium d. Some isotopes are more common than others i. most common isotope of hydrogen is protium

28. e. Radioisotopes: emit radiation and decay into other isotopes i. Tritium is an unstable isotope of hydrogen and decaysover time ii. Continue to decayuntil the isotope reaches a stable point

29. Protium f. Number of neutrons can be calculated i. write the mass number and atomic number of the isotope before thesymbolof the element

30. ii. Number of neutrons can be found by subtractingthe atomic number from the mass number iv. Example: uranium (U) -235 Mass number (A) = 235 - Atomic number (Z) = 92_ 143

31. D. Atomic Masses a. because working with tiny masses is difficult, atomic masses are usually expressed inunified mass units(u)or atomic mass unit, amu b. Example: Carbon-12 i. Isotope of carbon ii. Has 6 protons and 6 neutrons iii. Each individual proton and neutron has a mass of about 1.0 u

32. Section 3: Modern Atomic Theory pp. 128 - 132

33. Modern Models of the Atom a. In 1913, Bohr (Danish physicist), proposed that the energy of each electron was related to the electron’s path around the nucleus

34. i. Can only be in certain energy levels ii. Must gain energy to move to a higher energy level iii. Must lose energy to move to a lower energy level

35. b. Exact location of an electron cannot be determined i. Orbital: shaded region where there is a likelihood of finding an electron, the darker the shading, the better the chance of finding an electron

36. B. Electron Energy Levels a. The number of energy levels that are filled in an atom depends on the number of electrons

37. b. Valence electrons: electrons in the outer energy level of an atom i. Determine the chemical properties of an atom

38. c. Four types of orbitals i. s, p, d, and f

39. ii. s orbitals are the simplest 1. only one possible orientation in space 2. shaped like a sphere 3. lowest energy 4. can hold up to 2 electrons

40. iii. p orbital 1. shaped like a dumbbell 2. can be oriented in space in 1 of 3 ways 3. each p orbital can hold up to 2 electrons for a total of 6 electrons

41. iv. d and f orbitals 1. more complex 2. 5 possible d orbitals 3. 7 possible f orbitals 4. each orbital can hold a maximum of 2 electrons

42. d. Orbitals determine the total number of electrons that each level can hold

44. C. Electron Transitions a. electrons jump between energy levels when an atom gains or loses energy

45. b. Ground State: lowest state of energy of an electron c. Excited State: when an electron gains energy d. Gain energy by absorbing a particle of light (photon)