1. MODELS OF THE ATOM
CHM 130
GCC

2. Review - Rounding Round 399
Hopefully you said 400 not 4. What is wrong with 4???
The zeroes ARE important, they are place holders and 4 are NOT the same! If you had $400 in the bank but they said you had only $4 you’d be pissed off!
Round 2389 to 2 sig fig, now round to 1 sig fig
Answer: NOT 24
2000 NOT 2

3. 5.1 DALTON’S ATOMIC THEORY
1. An element is made of tiny, indestructible particles called atoms. (Not quite true – why?)

4. 2. All atoms of an element are identical and have the same properties
2. All atoms of an element are identical and have the same properties. (Not quite true – why?)

5. 3. Atoms of different elements combine to form compounds.+

6. 4. Compounds contain atoms in small whole number ratios. e. g
4. Compounds contain atoms in small whole number ratios. e.g. Each H2O molecule consists of one O and two H atoms, not ½ atoms or ¾ atoms.

7. 5. Atoms of 2 or more elements can combine to form different compounds.
E.g. C and O may form CO or CO2

8. 5.2 Thomson cathode ray experiment

9. Thompson “Plum Pudding Model” http://highered. mcgraw-hill
Atom is + charged
e-’s are distributed throughout atoms like raisins in plum pudding + -

10. 5. 3 Rutherford’s Scattering Experiment http://www. mhhe

11. Explanation of Scattering
Figure: 05-03
Title:
Explanation of Alpha-Scattering
Caption:
(a) Powerful alpha particles should pass through a homogeneous "plum-pudding" gold atom. (b) Since positively-charged alpha particles are deflected, Rutherford reasoned correctly that a gold atom has a dense, positively-charged nucleus.
Notes:
Rutheford's observations were consistent with the atomic model in (b), an atom with a core of positive charge.

12. Nuclear Model
1) The atom is mostly empty space with electrons moving around.
2) Each atom has a small, dense nucleus with the Protons & Neutrons.

13. Rutherford’s model atom (~10-8 cm diameter) nucleus+ -
nucleus
(~10-13 cm diameter)
atom (~10-8 cm diameter)
If nucleus = size of a small marble, then atom is the size of Cardinal’s stadium!

14. Subatomic Particles

15. 5.4 Atomic Notation Figure: 05-04-01UN Title: Atomic Notation Caption:
In atomic notation, the mass number, the atomic number, and the chemical symbol are shown.
Notes:
Given the atomic notation of an atom, you can determine the number of protons, neutrons, and electrons in the atom. The number of neutrons is the difference between the mass number and the atomic number.

16. ATOMIC NUMBER
Every atom of an element has the same # of protons
The # of protons defines an element
Carbon atoms ALWAYS have 6 protons

17. So how calculate # neutrons?
Mass number
Mass number = # protons + neutrons
So how calculate # neutrons?
# neutrons = mass - # protons

18. Isotopes
Isotopes of an element have the same atomic number (# pro), but a different mass number (# neu).
Ex: carbon-12, carbon-13 and carbon-14
How many protons do the above have? Neutrons? 6
6, 7, 8

19. Ex. 1: Write the atomic notation for potassium-40.
How many neutrons are there?
______________
40 – 19 = 21

20. Ex. 2: a. Write the atomic notation for bromine-81. b
Ex. 2: a. Write the atomic notation for bromine b. How many neutrons are there? _________________
= 46

22. 5.5 Atomic Mass
Masses of atoms are so small that we define the atomic mass unit (amu)
Mass of proton & neutron  1 amu.
Mass of electron is basically zero amu

23. Atomic mass
Atomic Mass in the P. Table is the weighted average of all atoms for that element in the world, so that is why it is NOT a whole number.

24. Natural isotopes of carbon:
The atomic mass reported for carbon (12.01 amu) is closer to carbon-12 since it is most abundant isotope for C. (There is a ton more C-12 than C-13.)

25. Example: Use the Periodic Table to determine the most abundant isotope: a. lithium-6 or lithium b. chlorine-35 or chlorine-37

26. 5.6 Light has two components: Wavelength () is the distance between peaks Frequency () is the number of wave cycles per second. (like a beat)

27. As wavelength , the frequency , and the energy 

28. Which wave has higher energy? Lower frequency?

29. Radiant Spectrum:

30. 5.7 In 1900, Max Planck proposed the controversial idea that energy was emitted in small bundles called quanta.
a particle of light energy is called a photon

31. Ball loses potential energy continuously as it rolls down a ramp.
Ball loses potential energy in quantized amounts as it bounces down a stairway.
Figure: 05-10
Title:
Stair Analogy for Quantum Principle
Caption:
A ball rolling down a ramp loses PE continuously. A ball rolling down stairs loses PE in specific units.
Notes:
This diagram introduces the concept of electrons existing in discrete energy levels, like steps on a flight of stairs.

32. 5.8 Bohr Model ~1913
Neils Bohr proposed that electrons orbit around the nucleus, occupying orbits with distinct energy levels Electrons are quantized!

33. Bohr model of the atom
The electrons orbit around the nucleus kinda like planets orbit around the sun but in 3D.
These orbits are called energy levels or shells.
Each orbit has a specific radius and energy, so a certain distance from the nucleus.

34. Bohr Model
The orbit closest to the nucleus is lowest in energy; the energy increases with distance from the nucleus. Proven by line spectra.

35. When the light from a heated element passes through a prism, a series of narrow lines is seen. These lines are the emission line spectrum.

36. Atomic Fingerprints
Each element produces a different emission line spectrum, so its own unique color.

37. Each element has it’s own energy levels that are unique.

38. Bohr theory explains 3 lines in H2 spectra.
Electrons gain energy from heat or electricity and jump to a higher energy level. These “excited” electrons ultimately lose energy and drop to lower energy levels, which causes light to be emitted.

39. 5.9 Each Energy Level Can Be Subdivided Into Sublevels.
sublevels: s, p, d, and f.

40. Each Level Has n Sublevels: 1st level has One Sublevel
Each Level Has n Sublevels: 1st level has One Sublevel 1s 2nd level has Two Sublevels 2s 2p 3rd level - Three Sublevels 3s 3p 3d 4th level - Four Sublevels 4s 4p 4d 4f This is depicted on next slide.

42. Orbitals are regions in space where there is a high probability of finding an electron.
One orbital can hold a maximum of 2 electrons.

43. Each sublevel contains a specific number of orbitals.
s has 1 orbital
p has 3
d has 5
f has 7
Orbitals are boxes on next slide

45. 5.10 Electron Configuration: Shorthand description of electrons by sublevel.
Sublevels are filled in order of increasing energy.
1s < 2s < 2p < 3s < 3p < 4s
You will do configurations for the 1st 20 elements.
Note the 3d sublevel is higher in energy than the 4s which is why we fill 4s first

46. Writing electron configurations
# of electrons?
Fill in sublevels to reach that #
Use superscript numbers to indicate number of e-'s in each sublevel.
Ex: C is 1s22s22p2 (cause 6 electrons)

48. Practice writing e- configurationsNa O Ca Cl
1s22s22p63s1
1s22s22p4
1s22s22p63s23p64s2
1s22s22p63s23p5

49. The Periodic Table actually is arranged by s, p, d, and f sublevels.

50. Figure: 06-06
Title:
Blocks of Elements
Caption:
The relationship between energy sublevels and the s, p, d, and f blocks of elements is shown.
Notes:
Each block of elements represents a specific set of sublevels and is designated by a distinct color. THe indicated sublevel is the highest energy subshell in the atom.

51. S Orbitals 1s 3s 2s

52. P Orbitals

53. Cool orbital pictures http://winter.group.shef.ac.uk/orbitron/

54. Ch. 5 Self Test p. 140
Try # 1-4, 6-7, 9, 12, 14-15, 17-18
Answers in Appendix J
Try to answer first, then check your answer!
Also work the problems in the online NOTES and worksheets.