1. Ch 3: Atoms Problem Set Ch 3: page 89-90 1-3, 7-9, 12-13, 17-20, 22-24
Chapter 3 problem set:
Problem Set Ch 3: page 89-90
1-3, 7-9, 12-13, 17-20, 22-24

2. 3.1 The Atom: From Idea to Theory
Historical Background- In approximately 400 BC, Democritus (Greek) coins the term
“atom” (means indivisible). Before that matter was thought to be one continuous piece - called the continuous theory of matter. Democritus creates the discontinuous theory of matter. His theory gets buried for thousands of years
18th century - experimental evidence appears to support the idea of atoms.

3. Law of Conservation of Mass – Antoine Lavosier (French) -1700’s
The number of each kind of atoms on the reactant side must equal the number of each kind of atoms on the product side
A + B + C —> ABC

4. Law of Multiple Proportions – John Dalton (English) - 1803
The mass of one element combines with masses of other elements simple in whole number ratios.
Water (H2O) is always: 11.2% H; % O
Sugar (C6H1206) is always: 42.1% C; 6.5% H; % O
Law of Multiple Proportions Video

5. Law of Multiple Proportions – John Dalton (English) - 1803
Ex1:
wt. of H wt. of O
H + O  H2O
H + 0  H2O
The ratio of O in H2O2 to O in H2O = 32/16 = 2:1 (small whole numbers

6. Dalton’s Atomic Theory

7. 3.2 The Structure of the Atom
Updating Atomic Theory
1870’s - English physicist William Crookes - studied the behavior of gases in vacuum tubes(Crookes tubes - forerunner of picture tubes in TVs).
Crookes’ theory was that some kind of radiation or particles were traveling from the cathode across the tube. He named them cathode rays .

8. 3.2 The Structure of the Atom
20 years later, J.J. Thomson (English) repeated those experiments and devised new ones. Cathode Ray Tube
Thomson used a variety of materials, so he figured cathode ray particles must be fundamental to all atoms discovery of the electron.

9. 3.2 The Structure of the Atom
Thomson and Milliken (oil drop experiment) worked together (their data, not
them) to discover the charge and mass of the electron

10. 3.2 The Structure of the Atom
Charge and Mass of the electron - Thomson and Milliken (oil drop experiment) worked together to discover the charge and mass of the electron Oil Drop
charge = x coulomb this is the smallest charge ever detected
mass = 9/109 x g this weight is pretty insignificant

11. 3.2 The Structure of the Atom
Gold Foil Experiment (Rutherford - New Zealand)
Nuclei are composed of ‘nucleons’: protons and neutrons
Alpha particles from Polonium (in the lead box) were released towards a thin sheet of gold foil. Most of the particles went through and were seen on the detector screen. 1 in 20,000 alpha particles bounced back.

12. Rutherford’s Conclusion
Concluded:
1 – the positive portion of the atom is in the middle
2 – most of the atom is empty
3 – most of the mass is in the middle
4 – electrons orbit the nucleus
Analogy: if an atom is the size of the Linc, then the nucleus is the size of a tennis ball floating in the middle of the stadium.

13. Table: Subatomic particles important in chemistry.

14. Table: Subatomic particles important in chemistry.
Nuclear Forces Forces in Atoms

15. 3.3 Weighing and Counting Atoms
We look to the periodic table to give us information about the number of particles are in atoms and also to help us count atoms in a sample.
Counting nucleons
Atomic Number (Z) Atomic #
Number of protons in the nucleus
Uniquely labels each element
Mass Number (M) Mass #
Number of protons + neutrons in the nucleus

16. Counting nucleons

17. Counting electrons Atoms Ions Same number of electrons and protons
Ionic charge (q) = #protons - #electrons
Positive ions are cations
Negative ions are anions

19. Review of formulas
atomic # (Z) - (always a whole number, smaller number on the periodic table) = # of protons in the nucleus - also indicates the # of electrons if the element is not charged
atomic mass – the average mass of all of the isotopes of an element – is a number with a decimal – is always the larger number on the periodic table.
mass number (A) - sum of the protons and neutrons in a nucleus
this number is rounded from atomic mass due to the fact that there are isotopes
# neutrons = A - Z example - # of neutrons in Li = = rounds to 4
Ion – a charged atom. Atoms become charged by gaining electrons (become a negative charge) or losing electrons (become a positive charge)

20. Lots of Practice!!! p+ e- n° Atomic # = (# of p+) Mass # = (p+ + n0) C6 12 Ca 20 40 U 92
146
238 Cl 17 18 35 Mg 24
14C 8 14
S-2 16 32
Na+1 11 10 23

21. Isotopes Isotopes Isotope Video
Two atoms of the same element (same # of p+) but with different masses (different # of n0)

22. Average Atomic Mass (“weighted average”)
Definition - The average weight of the natural isotopes of an element in their natural abundance.
History lesson - originally H was the basis of all atomic masses and was given the mass of Later, chemists changed the standard to oxygen being (which left H = 1.008). In 1961, chemists agreed that 12C is the standard upon which all other masses are based.
1/12 of the mass of 1 atom of 12C = 1 amu

23. Carbon consists of two isotopes: 98. 90% is C-12 (12. 0000 amu)
Carbon consists of two isotopes: 98.90% is C-12 ( amu). The rest is C-13 ( amu). Calculate the average atomic mass of carbon to 5 significant figures.
(.9890)( )+(.0110)( )=x
=12.011

24. Ex1: Chlorine consists of two natural isotopes, 35Cl (34. 96885) at 75
Ex1: Chlorine consists of two natural isotopes, 35Cl ( ) at 75.53% abundance and 37Cl ( ) at 24.47% abundance. Calculate the average atomic mass of Chlorine.
(.7553)( )+(.2447)( )=x
=35.46
Ex2: Antimony consists of two natural isotopes 57.25% is 121Sb ( ). Calculate the % and mass of the other isotope if the average atomic mass is
(.5725)( )+(.4275)(x) =121.8
x=121.8
x= x=123.0

25. The Mole, Avogadro’s number and Molar Mass
The Mole Mole Video
Atoms are tiny, so we count them in “bunches”.
A mole is a “bunch of atoms”.
The Mole (definition) -The amount of a compound or element that contains 6.02 x particles of that substance. Avogadro
1 mole = 1 gram formula mass = x particles

26. Molar Mass - the sum of the atomic masses of all atoms in a formula
Round to the nearest tenth! (measured in amu or grams)
ex - H2 H2O Ca(OH)2
2.0g 18.0 g 74.1 g

27. Official names may also be:
Molar mass is a term that can be used for atoms, molecules (covalent compounds or elements) and formula units (ionic compounds)
Official names may also be:
Formula mass (ionic compounds)
Molecular mass (covalent compounds and diatomic elements)
Atomic weight, Atomic mass, grams formula weight, etc.
Molar Mass

28. Examples: 1 mole Na = 6.02 x 10 23atoms = 23.0g
1 mole O2 = x molecules=32.0g
1 mole HCl = x molecules = 36.5 g
1 mole NaCl =6.02 x formula units = 58.5 g

29. Mole Relationships 1 mole 6.02 x 1023 atom/molecule P-Table for gram
22.4 liter (at STP