1. Chapter 3 Atoms, Molecules, and Ions

2. Law of Definite Proportions
Also known as the Law of Constant Composition
Two samples of a given compound are made of the same elements in exactly the same proportions by mass, regardless of the size or sources of the samples
NaCl % Na Sodium % Cl Chlorine
H2O 11.1% H Hydrogen 88.9% O Oxygen

3. The Law of Conservation of Mass
The mass of the reacting system does not change
Take the following example:
Table salt
If I react 23 g of sodium
With 35.5 g of chlorine
I get 58.5 g of sodium chloride
Water
If I react 2 g of hydrogen
With 16 g of oxygen
I get 18 g of water

4. Law of Multiple Proportions
When two elements combine to form more than one compound, the mass ratios of the two compounds are related by small whole numbers
Examples
Hydrogen Peroxide (H2O2)
2 atoms of hydrogen react with two atoms of oxygen
Water (H2O)
? Atoms of hydrogen react with ? Atoms of oxygen.

5. Atomic Theory of Matter
John Dalton
All Matter is composed of extremely small particles called atoms. (cannot be created/subdivided or destroyed)
All atoms of a given element are identical to one another in chemical and physical properties
The atoms of one element are different from the atoms of all other elements in physical and chemical properties.
Atoms of different elements combine in simple whole number ratios to form compounds
In chemical reactions, atoms are combined, separated, or rearranged but never created, destroyed, or changed.
Figure 2.1 John Dalton ( )

6. The Electron – Discovered by J.J. Thompson
Streams of negatively charged particles were found to emanate from cathode tubes.
J. J. Thompson is credited with their discovery (1897).
Thompson measured the charge/mass ratio of the electron to be 1.76  108 coulombs/g.
Figure 2.4

7. The Atom, circa 1900: “Plum pudding” model, put forward by Thompson.
Positive sphere of matter with negative electrons imbedded in it.
Figure 2.9

8. Discovery of the Nucleus
Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.
The Experiment
Figure 2.10

9. The Nuclear Atom
Since some particles were deflected at large angles, Thompson’s model could not be correct.
Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom.
Most of the volume of the atom is empty space.
Figure 2.11

10. Other Subatomic Particles
Protons were discovered by Rutherford in 1919.
Neutrons were discovered by James Chadwick in 1932.

11. Subatomic Particles
Protons and electrons are the only particles that have a charge.
Protons and neutrons have essentially the same mass.
The mass of an electron is so small we ignore it.
Table 2.1
Location
Charge
Mass unit
Proton
Nucleus +1 1
Neutron
Electron
Orbitals -1

12. Symbols of Elements
Elements are symbolized by one or two letters. 12

13. Atomic Number
All atoms of the same element have the same number of protons:
The atomic number (Z) 13

14. Atomic Mass
The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom. 14

15. Isotopes:
Atoms of the same element with different masses.
Isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C 15

16. Electron Configurations
The way electrons are arranged in atoms.
Aufbau principle- electrons enter the lowest energy first.
This causes difficulties because of the overlap of orbitals of different energies.
Pauli Exclusion Principle- at most 2 electrons per orbital - different spins 16

17. Electron Configuration
Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to . 17

18. Atomic Orbital Model for Electrons

19. S orbitals 3s 2s 1s 1 s orbital for every energy level
Spherical shaped
Each s orbital can hold 2 electrons
Called the 1s, 2s, 3s, etc.. orbitals. 3s 2s 1s 19

20. P orbitals Start at the second energy level 3 different directions
3 different shapes
Each can hold 2 electrons 20

21. P Orbitals

22. D orbitals Start at the third energy level 5 different shapes
Each can hold 2 electrons 22

23. Summary # of shapes Max electrons Starts at energy level s 1 2 1 p 3 6d 5 10 3 7 14 4 f

24. Number of Electrons that can be held in each orbital
s = 2 e-
p = 6 e-
d = 10 e-
f = 14 e-

25. By Energy Level First Energy Level only s orbital only 2 electrons 1s2
Second Energy Level
s and p orbitals are available
2 in s, 6 in p
2s22p6
8 total electrons 25

26. By Energy Level Third energy level s, p, and d orbitals
2 in s, 6 in p, and 10 in d
3s23p63d10
18 total electrons
Fourth energy level
s,p,d, and f orbitals
2 in s, 6 in p, 10 in d, ahd 14 in f
4s24p64d104f14
32 total electrons 26

27. By Energy Level
Any more than the fourth and not all the orbitals will fill up.
You simply run out of electrons
The orbitals do not fill up in a neat order.
The energy levels overlap
Lowest energy fill first. 27

28. Increasing energy 7p 6d 7s 6p 5f 5d 6s 5p 4f 5s 4d 4p 3d 4s 3p 3s 2p28

29. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
The next electrons go into the 2s orbital
only 11 more 29

30. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The next electrons go into the 2p orbital
only 5 more 30

31. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The next electrons go into the 3s orbital
only 3 more 31

32. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s
The last three electrons go into the 3p orbitals.
They each go into seperate shapes
3 upaired electrons
1s22s22p63s23p3 32

33. The easy way to remember
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f 33

34. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2
4 electrons 34

35. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2
12 electrons 35

36. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2
20 electrons 36

37. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
38 electrons 37

38. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
56 electrons 38

39. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2
88 electrons 39

40. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6
108 electrons 40

41. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d4f 5f 41

42. Determining Quantum Numbers n,l,m,s
n – Principal Quantum Number – same as energy level
l – Angular Quantum Number – determined by orbital
s=0 p=1 d=2 f=3

43. Determining Quantum Numbers n,l,m,s
m – Magnetic Quantum number– for s orbital = 0. Others determined be which one. Example is for the p & d orbitals. p d
s – Spin Quantum Number
either + ½ ( ↑) or – ½ (↓) -1 +1 -2
- 1
+ 1` +2

44. Periodic Table: A systematic catalog of elements.
Elements are arranged in order of atomic number.

45. Periodicity
When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.

46. Periodic Table The rows on the periodic chart are periods.
Columns are groups.
Elements in the same group have similar chemical properties.

47. Groups
These five groups are known by their names.

48. Periodic Table
Nonmetals are on the right side of the periodic table (with the exception of H).

49. Periodic Table
Metalloids border the stair-step line (with the exception of Al and Po).

50. Periodic Table
Metals are on the left side of the chart.

51. The average of the atomic masses of the naturally occurring isotopes
Atomic Mass Unit (amu)
The average of the atomic masses of the naturally occurring isotopes

52. SI unit for the amount of a substance
Mole (mol)
SI unit for the amount of a substance
The number of atoms or particles in exactly 12 grams of carbon-12

53. Molar Mass The mass in grams of one mole of the element
Used to convert between moles and grams
Unit: g/mol
Note: The mass in grams of 1 mol of an element is numerically equal to the element’s atomic mass from the periodic table in atomic mass units

54. Example: 10 grams of Cu = ? mol of Cu
Atomic mass of Cu: amu Molar mass of Cu: g/mol
10 g x 1 mol = 0.16 mol g

55. Example 2: What is the mass in grams of 3. 50 mol of Lithium
Example 2: What is the mass in grams of 3.50 mol of Lithium? Molar mass of Li = 6.94 g/mol Amount of Li = 3.50 mol
Calculate: mol x grams Li = grams Li mol Li

56. Summary Convert from moles to mass: ____ mol x ___ grams =
Summary Convert from moles to mass: ____ mol x ___ grams = ? Grams mol Convert from mass to moles: ____ grams x mol = ? Mol ___ grams

57. The number of atoms or molecules in 1 mol of a substance
Avogadro’s Number
The number of atoms or molecules in 1 mol of a substance
1 mole of pure substance contains x 1023 particles

58. Example: Determine the number of atoms in 0.30 mol of fluorine atoms.
Amount of F = 0.30 mol
Number of atoms = ?
0.30 mol F x x 1023 F atoms
1 mol F
= 1.8 x 1023 Fluorine atoms

59. Summary: Moles to atoms: Atoms to moles:
____ mol x x 1023 atoms = ? atoms
1 mol
Atoms to moles:
___ atoms x mol = ? Moles
6.022 x 1023 atoms