1. Ch 4: Atoms

3. Scientists (key experiments and contributions)
Democritus
Lavosier
Dalton
Crookes
Thomson
Miliken
Rutherford
Counting P,N,E
Atoms, Ions, Isotope
Cation vs anion
Define Ion and Isotope
Average Atomic Mass Problems

4. 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.

5. 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 + 2B + C —> AB2C

6. 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

7. Dalton’s Atomic Theory
Everything is made of atoms
2. Atoms of the same element are identical (NOT TRUE)
3. Atoms can not be broken down, created or destroyed. (NOT TRUE)
4. Atoms combine in simple whole number ratios to form chemical compounds
5. A chemical reaction is the combining, separation, or rearrangement of atoms.

8. 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 .

9. 3.2 The Structure of the Atom (video)
20 years later, J.J. Thomson (English) repeated those experiments and devised new ones.
Thomson used a variety of materials, so he figured cathode ray particles must be fundamental to all atoms discovery of the electron.
Plum Pudding Model

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
charge = x coulomb this is the smallest charge ever detected
mass = 9.11 x g this weight is pretty insignificant

11. 3. 2 The Structure of the Atom https://www. youtube. com/watch
Gold Foil Experiment (Rutherford - New Zealand)
Nuclei are composed of ‘nucleons’: protons and neutrons

12. A) Oil Drop B) Found Nucleus C) Law of conservation of mass D) Plum Pudding Model E) Coined the term cathode ray F) Term atom G) Law of multiple proportions H) Cathode Ray Tube experiments I) Gold Foil Experiment J) Charge and Mass of electron
1) Lavosier 2) Democritus 3) Milliken 4) Rutherford 5) Dalton 6) Thomson 7) Crookes

13. Table: Subatomic particles important in chemistry.

14. 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.
Atomic Number (Z)
-Number of protons in the nucleus
- Uniquely labels each element
Mass Number (M)
- Number of protons + neutrons in the nucleus

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

18. 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)

19. Practice- How many Protons?23 16 7

20. Practice- How many Electrons?23 16 7

21. Practice- How many Neutrons?23 16 7

22. Lets try a few together

23. Lots of Practice!!! p+ e- n° Atomic # = (# of p+) Mass # = (p+ + n0) CCa U Cl Mg
14C
S-2
Na+1

24. Protons
Neutrons
Electrons Sr
Sr +2
O -2 H+
B +3

25. Protons
Neutrons
Electrons 38 50 36 8 10 1 5 6 2 Sr
Sr +2
O -2 H+
B +3

26. Complete Practice Problems

27. Isotopes
Isotopes
Two atoms of the same element (same # of p+) but with different weights (different # of n0)
H-1, H-2,H C-12, C-13, C U-235, U-238

28. How do you determine your grades in this class?
60% tests: 90 40% other: 100

29. Average Atomic Mass (“weighted average”)
Definition - The average weight of the natural isotopes of an element in their natural abundance.

30. 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

31. 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
=35.46amu
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
.4275x=
=123.0amu

32. Practice Worksheet Problems

33. Beanium Lab

34. Scientists (key experiments and contributions)
Democritus
Lavosier
Dalton
Crookes
Thomson
Miliken
Rutherford
Counting P,N,E
Atoms, Ions, Isotope
Cation vs anion
Define Ion and Isotope
Average Atomic Mass Problems