1. Atoms: The Building Blocks of Matter
Ch. 3

2. Early Atomic Theory Democritus: coined the word atom (“indivisible”)
Thought nature’s particle was an atom
Aristotle: didn’t believe atoms existed
Thought all matter was continuous
Neither theory was supported by scientific evidence.
Evidence was not gathered until the 18th century.

3. Early Atomic Theory
Late 1700s: all scientists believed in the concept of an element
Elements combined to form compounds.
Chemical reaction: combination of a substance or substances into one or more new substances
Improved equipment helped chemists to more accurately measure experiments

4. Basic Laws of Chemistry
Law of conservation of mass: mass is neither created nor destroyed during ordinary chemical rxns or physical changes
Law of definite proportions: a chemical compound contains the same elements in exactly the same proportions by mass regardless of sample size
Law of multiple proportions: If two or more different compounds are composed of the same two elements, then the ratio of masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers

6. Dalton’s Atomic Theory
English Schoolteacher (1808)
Proposed an explanation for the three laws described above
Theory:
All matter is composed of atoms.
Atoms of a given element are identical.
Atoms cannot be subdivided, created, or destroyed.
Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
In chemical reactions, atoms are combined, separated, or rearranged.

7. How does Dalton’s law explain chemical laws?
Law of conservation of mass: mass is not changed – it is rearranged, separated, or combined
Law of definite proportions: a given chemical compound is always composed of the same combination of atoms
Law of multiple proportions: ratio of different masses of the same compound

8. Modern Atomic Theory
Not all Dalton’s statements have been proven to be correct.
We can divide atoms.
Atoms of a given element can have different masses (isotopes).
Important remaining concepts:
All matter is composed of atoms
Atoms of any one element differ in properties from atoms of another element

9. Atomic Structure
Atom: smallest particle of an element that retains the chemical properties of that element
All atoms have two regions:
Nucleus:
Proton: positively charged particle
Neutron: neutral particle
Electrons: negatively charged particles

10. Discovery of the Electron
Resulted from investigations between electricity and matter (late 1800s)
Many experiments were conducted using cathode-ray tubes: pass electric current through a tube of gas
Noticed that when electricity was passed through tube, opposite end glowed.
They thought the glow was from a stream of particles (a cathode ray).
They found these rays were deflected by negatively charged objects

12. Charge/Mass of Electron
Robert A. Millikan: found the charge of an electron
Mass of an electron: x kg.
Two inferences made on atomic structure:
Because atoms are electrically neutral, they must contain a positive charge to balance the negative electrons.
Because electrons have so much less mass than atoms, atoms must contain other particles that account for most of their mass.

15. Atomic Number All atoms are composed of the same basic particles
Atoms of different elements have different numbers of protons
Atomic number: number of protons of each atom in that element

16. Isotopes Some elements have different numbers of neutrons.
Example: Hydrogen
Protium: 1 proton, 1 electron
Deuterium: 1 proton, 1 neutron, 1 electron
Tritium: 1 proton, 2 neutrons, 1 electron
Isotopes: atoms of the same element that have different masses
Mass number = number of protons + number of neutrons

18. Designating Isotopes
Most isotopes don’t have unique names like hydrogen does
Usually identified by their mass number
Two ways to identify them
Hyphen notation
Hydrogen-3 = Tritium
Number at the end represents the mass number
Nuclear symbol: 168O = oxygen
Top number is the mass number, bottom number is the atomic number.

19. Sample Problem A
How many protons, electrons, and neutrons are there in an atom of chlorine-37?

20. Atomic Mass
Masses of atoms are very small…kind of cumbersome to use grams
Atomic mass unit: 1/12 the mass of a carbon-12 atom
Example: Hydrogen – 1.08 amu (located on periodic table)
Average atomic mass: weighted average of the atomic masses of the isotopes of an element

21. Average Atomic Mass Box with two sizes of marbles: 100 marbles
25% = 2.00g each
75% = 3.00g each
100 marbles
25 marbles x 2.00 g = 50 g
75 marbles x 3.00 g = 225 g
= 275 g
275/100 = 2.75 g (average mass of the marbles)

22. Calculating Average Atomic Mass
Copper
69.15% copper-63 (amu = )
30.85% copper 65 (amu = )
Average atomic mass
( x ) + ( X ) =
63.55 amu
(always round atomic mass to 2 decimal places before using in any calculation)

23. Sample Problem (p.82 Table 4)
Hydrogen, carbon, oxygen

24. The Mole
Definition: the amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12.
A counting unit – just like a dozen is.
Avogadro’s number: the number of particles in exactly 1 mole of a pure substance (6.022 x 1023)
1 mole = x 1023 atoms

25. Molar Mass The mass of one mole of pure substance Units: g/mol
Numerically equivalent to atomic mass
Look on the number underneath each element of the periodic table
CONVERSIONS! Remember to put whatever you want to find on top.

26. Sample Problem B
What is the mass in grams of 3.5 mol of the element copper (Cu)?

27. Sample Problem C
A chemist produced 11.9 g of aluminum, Al. How many moles of aluminum were produced?

28. Sample Problem D
How many moles of silver, Ag, are in 3.01 x 1023 atoms of silver?

29. Sample Problem E
What is the mass in grams of 1.20 x 108 atoms of copper, Cu?