1. Chapter #3 Atoms: The Building Blocks of Matter Chapter 3 ATOMS:
West Valley High School
General Chemistry
Mr. Mata
Chapter #3
Atoms: The Building Blocks of Matter

2. Standard 2a
To compare and contrast the atomic number and atomic mass.

3. Essential Question
Describe how atomic theory developed to give us our modern view of the atom.

4. 3-1 Early Atomic Theory Atoms – smallest particle of matter.
Democritus (400 B.C) stated world was made of atoms. (atomos = “indivisible”).
Antoine Lavoisier (1800’s) discovered mass didn’t change after chemical rxn.
Proposed “matter can be changed, but it cannot be created or destroyed “
(Law of Conservation of Mass).

5. Dalton’s Atomic Theory
All matter composed of atoms.
Atoms of same element identical in size, mass, properties; atoms of diff elements diff. in size, mass, properties.
Atoms can’t be subdivided, created, destroyed.
Atoms of different elements combine in whole-number ratios to form chemical compounds.
In chemical reactions, atoms are combined, separated, or rearranged.

6. Modern Atomic Theory All matter is composed of atoms.
Atoms of any one element differ in properties from atoms of another element.
Element’s average mass unique to element.
Atoms cannot be subdivided, created, or destroyed in ordinary chemical rxns.
Changes CAN occur in nuclear rxns!

7. Atomic Structure
Atom- smallest particle of element that retains chemical properties of element.
Nucleus- positively charged, dense central portion of the atom; contains nearly all mass (~ 99.7%).

8. Subatomic Particles Electrons e-
Negative charged particles. Found in electron clouds. Thomson (1897)
Protons p+
Positively charged particles. Found in the nucleus. Rutherford (1918)
Neutrons N
No charge. Found in the nucleus. Chadwick (1932)

9. The Atomic Scale
Most of the mass of the atom is in the nucleus (protons and neutrons)
Electrons are found outside of the nucleus (the electron cloud).
Most volume of atom=empty space.

10. Famous Scientists Scientist Experiment Name Conclusion JJ Thomson
Cathode Ray
Electron (-)
Robert Millikan
Oil Drop
Mass of a single electron.
Ernest Rutherford
______________
James Chadwick
Gold foil
_______________
Beryllium Detection
Proton (+)
_____________
Neutron (0)

11. Discovery of the Electron
J.J. Thomson (1897) used a cathode ray tube to deduce the presence of a negatively charged particle.
Discovered the – electron particle.

12. (1906)Thomson awarded the Nobel Prize in chemistry for discovery of the electron.
The atom could be broken down into smaller particles.

13. Thomson’s Atomic Model
Thomson believed electrons were like plums in a + charged “pudding”.
He called it the “plum pudding” model.

14. Rutherford’s Gold Foil Experiment
Alpha particles are helium nuclei.
Particles fired at a thin sheet of gold foil.
Particle hits on screen (film) are detected.

15. Rutherford’s Findings
Most particles passed right through screen.
Few particles deflected.
VERY FEW were greatly deflected.
Conclusions:
The nucleus is small.
The nucleus is dense.
The nucleus is + charged.

16. Section 3-3
Atomic number: number of protons in the nucleus of atom.
# p(+) = # e(-) 6 C
Carbon
12.011

17. Atomic Mass
number of protons & neutrons in the nucleus.
Atomic Mass=protons + neutrons
Atomic Mass C =
Mass number = rounded atomic mass
Mass Number C = 12

18. 235 92 Nuclear Symbols Mass number (p+ + n) Atomic Mass (p+ + n)
Element symbol
Atomic number (# of p+) U
235 92

19. Hyphen Notation Sodium-23 (23 is the atomic mass)
Sooo… (atomic #) = 12 for the # of neutrons.
Atomic number of 11 is the # of protons (11) and electrons(11).

20. Molar Relationships
1 mole = 6.02 x 1023 atoms, molecules, or particles
Molar mass = mass of 1 mole of substance(Hint: Add up atomic masses)
Ex: Find the molar mass of:
H2O = (H=1(2), 0=16), 18 grams
CO2 = (C=12, O=16(2)), 44 grams
H2SO4=(H=1(2),S=32,O=16(4)), 98 g

21. Molar Masses Find molar masses: NaCl = 58 grams = 1 mole
CaCl2 = 110 grams = 1 mole
Fe(OH)2 = 90 grams = 1 mole
H3PO4 = 98 grams = 1 mole
C6H12O6 = 180 grams = 1 mole

22. Calculations: Converting moles to grams
Given # of mole X ? g (look at periodic table)= g
1 mole
How many grams of lithium are in 3.50 moles of lithium?
3.50 mole Li X 7 g = g Li
1 mol

23. Your turn… How many grams of carbon are in 8.25 moles of carbon?
8.25 mole C X 12 g = 99.0 g C
1 mol

24. Your turn… How many grams of uranium are in 21.5 moles of uranium?
21.5 mole U X 238 g = g U
1 mol

25. Calculations: Converting grams to moles
Given # of g X 1 mol = mol
g (look at periodic table)
How many moles of lithium are in 18.2 grams of lithium?
18.2 g Li X 1 mol Li = 2.6 mol Li
7 g

26. Your turn… How many moles of gold are in 150 grams of gold?
150 g Au X 1 mol Au = 0.8 mol Au
197 g

27. Your turn… How many moles of H2O are in 120 grams of water?
120 g H2O X 1 mol H2O= 6.7 mol H2O
18 g

28. I didn’t discover it. Its just named after me!
Avogadro’s Number
Number of particles in exactly one mole of a pure substance.
6.02 x 1023 is called “Avogadro’s Number”.
Named in honor of the Italian chemist Amadeo Avogadro ( ).
I didn’t discover it. Its just named after me!

29. Important Skill
To find electrons, protons, and neutrons for an element:
40 Ca  atomic mass (#p + #n)
 atomic number (#p = #e)
# electrons = 20 (atomic number)
# protons = 20 (atomic number)
# neutrons = 20 (atomic mass–atomic #)
40 – 20 = 20 neutrons

30. Chapter 3 SUTW Prompt
Describe the contributions of Thomson, Millikan, & Rutherford to our understanding of the atom.
Complete an sentence paragraph using the SUTW paragraph format. Hilite using green, yellow, and pink.
Due Date: Monday, September 11, (beginning of class).