1. Chapter 2 Atoms, Molecules and Ions

2. Problems from Chapter 2
Pages
1 – 6, 9-12, 14, 16-19, 22,
24-34, 36-46, 48, 50, 60

3. EH Assignment
Due
Unit 3 Chemical Nomenclature
Minimum score
Sec 1 Names/Symbols of Elements 90
Sec 2 Molecules and Ions
Sec 3 Cations and Anions 85
Sec 4 Ionic Compounds
Sec 5 Acids, Bases, and Salts 80
Sec 6 Elements, Compound and Mixtures

4. EH Assignment
Due
Unit 9 Atomic Structure
Minimum score
Sec 1 Elementary Particles and Isotopes 90
Test on Chapters 1 and 2

5. Dalton’s Atomic Theory (1808)
Elements are composed of extremely small particles called atoms. All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements.
Compounds are composed of atoms of more than one element. The relative number of atoms of each element in a given compound is always the same.
Chemical reactions only involve the rearrangement of atoms. Atoms are not created or destroyed in chemical reactions.
2.1

6. Dalton developed the atomic theory to explain
The Law of Conservation of Mass
The Law of Definite Proportions
The Law of Multiple Proportions
Laws are summaries of experimental observations.
Theories are models devised to explain laws.

8. Laws of Mass Conservation & Definite Composition
Law of Mass conservation: The total mass of substances
does not change during a chemical reaction.
Law of Definite ( or constant ) composition: No matter
what its source, a particular chemical
compound is composed of the same elements
in the same proportions by mass.

9. Law of Conservation of Mass
16 X
8 Y +
8 X2Y
Law of Conservation of Mass
2.1

10. LAW OF MULTIPLE PROPORTIONS
If two elements can combine to form more than one compound, the masses of one element that combine with with a fixed mass of the other element are in the ratio of small whole numbers.
Carbon monoxide g C to 1.33 g O
Carbon dioxide g C to g O

11. Law of Multiple Proportions2
Law of Multiple Proportions
2.1

12. THE STRUCTURE OF THE ATOM
Dalton thought that the atom was indivisible, but by the end of the 19th century there was evidence that the atom was composed of smaller particles.
The electron was discovered in the 1890’s by J. J. Thomson using a Crooks tube (cathode ray tube).
He was able to measure the ratio of the electric charge to mass of the electron to be x 108 C/g.
Later an American physicist, Millikan, measured the charge to be x C.

13. Fig. 2.4a

14. Fig. 2.4b

15. J.J. Thomson, measured mass/charge of e- (1906 Nobel Prize in Physics)
2.2

16. Thomson’s charge/mass of e- = -1.76 x 108 C/g
Measured mass of e-
(1923 Nobel Prize in Physics)
e- charge = x C
Thomson’s charge/mass of e- = x 108 C/g
e- mass = 9.10 x g
2.2

17. The mass of the electron was much
less than the mass of the lightest atom.
This led to Thomson’s “Plum Pudding” model of the atom.

18. 2.2

19. (Uranium compound)
2.2

20. Rutherford Experiment
Bombarded a thin gold foil with high-energy alpha particles from radium.
Most alpha particles went through unaffected.
About 1 in 10,000 deflected through a large angle.
Lead to nuclear model of atom.

21. (1908 Nobel Prize in Chemistry)
particle velocity ~ 1.4 x 107 m/s
(~5% speed of light)
atoms positive charge and most of its mass
are concentrated in the nucleus
2. light electrons are in the outer part of the atom.
2.2

22. Rutherford’s Model of the Atom If the atom is the Houston Astrodome
atomic radius ~ 100 pm = 1 x m
nuclear radius ~ 5 x 10-3 pm = 5 x m
If the atom is the Houston Astrodome
Then the nucleus is a marble on the 50 yard line
2.2

23. By the 1930’s it was known that the nucleus contained two subatomic particles - the proton and the neutron. The nucleus has a volume which is only a tiny fraction of the atom.
The electrons occupy the outer part of the atom.

24. Subatomic Particles (Table 2.1)
mass p = mass n = 1840 x mass e-
2.2

25. X H H (D) H (T) U Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei
Mass Number X A Z
Element Symbol
Atomic Number H 1
H (D) 2
H (T) 3 U
235 92
238
2.3

26. 2.3

27. Do You Understand Isotopes?
How many protons, neutrons, and electrons are in C 14 6 ?
6 protons, 8 (14 - 6) neutrons, 6 electrons
How many protons, neutrons, and electrons are in C 11 6 ?
6 protons, 5 (11 - 6) neutrons, 6 electrons
2.3

28. THE PERIODIC TABLE
Elements with similar chemical and physical properties are grouped together in vertical columns called GROUPS or FAMILIES. Rows are called PERIODS.
Group 1 (1A) - alkali metals
Group 2 (2A) - alkaline earth metals
Group 17 (7A) - halogens
Group 18 (8A) - noble gases
Transition elements (3-12)
metals, nonmetals, metalloids

29. Alkali Earth Metal
Noble Gas
Halogen
Alkali Metal
Period
Group
2.4

30. A diatomic molecule contains only two atoms
A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical bonds H2
H2O
NH3
CH4
A diatomic molecule contains only two atoms
H2, N2, O2, F2, Cl2, Br2, I2, HCl, CO
A polyatomic molecule contains more than two atoms
O3, H2O, NH3, CH4
2.5

31. cation – ion with a positive charge
An ion is an atom, or group of atoms, that has a net positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation. Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl-
17 protons
18 electrons Cl
17 protons
17 electrons
2.5

32. A monatomic ion contains only one atom
Na+, Cl-, Ca2+, O2-, Al3+, N3-
A polyatomic ion contains more than one atom
OH-, CN-, NH4+, NO3-
2.5

33. How do we know how many electrons are
lost or gained?
For elements near either side of the periodic table the atoms tend to form
ions with the same number of electrons
as the nearest noble gas.
What kind of ions do these elements form?
I, Ba, Al, S, Rb, Fe

34. A Polyatomic Ion
Fig. 2.22

35. How many protons and electrons are in
Do You Understand Ions?
How many protons and electrons are in Al 27 13 ? 3+
13 protons, 10 (13 – 3) electrons
How many protons and electrons are in Se 78 34 2- ?
34 protons, 36 (34 + 2) electrons
2.5

36. All metals form cations.
2.5

37. 2.6

38. TA p49

39. TA p49

40. An empirical formula shows the simplest
A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance
An empirical formula shows the simplest
whole-number ratio of the atoms in a substance
H2O
molecular
empirical
H2O
C6H12O6
CH2O O3 O
N2H4
NH2
2.6

41. The ionic compound NaCl
ionic compounds consist of a combination of cations and an anions
they only have empirical formulas, no molecular formula
the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero
The ionic compound NaCl
2.6

42. Formula of Ionic Compounds
2 x +3 = +6
3 x -2 = -6
Al2O3
Al3+
O2-
1 x +2 = +2
2 x -1 = -2
CaBr2
Ca2+
Br-
1 x +2 = +2
1 x -2 = -2
Na2CO3
Na+
CO32-
2.6

43. Some Polyatomic Ions (Table 2.3)
2.7

45. For our test you need to know:
Ammonium
Hydroxide
Nitrate
Sulfate
Carbonate Bicarbonate
Phosphate

46. This is an ionic compound.
This is the ammonium cation.
How can you tell if a compound is molecular
or ionic?

47. Which of these are ionic and which are molecular?
SiCl2, LiF, BaCl2, B2H6, KCl, C2H4,
NH4NO3

48. Chemical Nomenclature
Ionic Compounds (salts)
often a metal + nonmetal (or polyatomic ion)
anion (nonmetal), add “ide” to element name
BaCl2
barium chloride
K2O
potassium oxide
Mg(OH)2
magnesium hydroxide
KNO3
potassium nitrate
2.7

49. Transition metal ionic compounds
indicate charge on metal with Roman numerals
FeCl2
iron(II) chloride
2 Cl- -2 so Fe is +2
FeCl3
3 Cl- -3 so Fe is +3
iron(III) chloride
Cr2S3
3 S-2 -6 so Cr is +3 (6/2)
chromium(III) sulfide
2.7

50. TA p45

51. Name these:
LiF, MgCl2, NaOH, Al2(SO4)3, NH4NO3,
FeO
Write formulas for:
potassium carbonate, ammonium sulfide
cobalt(II) bromide, aluminum phosphate

52. Molecular compounds (binary)
nonmetals or nonmetals + metalloids
common names
H2O, NH3, CH4
element further left in periodic table is 1st
element closest to bottom of group is 1st
if more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom
last element ends in ide
2.7

53. Molecular Compounds HI hydrogen iodide NF3 nitrogen trifluoride SO2
sulfur dioxide
N2Cl4
dinitrogen tetrachloride
TOXIC!
NO2
nitrogen dioxide
N2O
dinitrogen monoxide
Laughing Gas
2.7

55. Practice exercises 2. 6 and 2
Practice exercises 2.6 and 2.7 Name NBr3 and Cl2O7 Write formulas for sulfur tetrafluoride dinitrogen pentoxide

56. 2.7

57. An acid can be defined as a substance that yields
hydrogen ions (H+) when dissolved in water.
HCl
Pure substance, hydrogen chloride
Dissolved in water (H+ Cl-), hydrochloric acid
An oxoacid is an acid that contains hydrogen, oxygen, and another element.
HNO3
nitric acid
H2CO3
carbonic acid
H2SO4
sulfuric acid
2.7

59. 2.7

61. If IO3- is the iodate ion name IO2-, HIO3, and HIO2

62. A base can be defined as a substance that yields
hydroxide ions (OH-) when dissolved in water.
NaOH
sodium hydroxide
KOH
potassium hydroxide
Ba(OH)2
barium hydroxide
2.7