1. Atoms: The building blocks of matter
Chapter 3

2. 3.1 The Atom
Democritus
First to support the idea of the atom as the smallest particle of matter
Circa 400 BC
Competing with other philosophers of the time as well
Idea falls flat because he doesn’t have any proof to support his claim

3. 3.1 The Atom John Dalton (1808) Atomic Theory
All matter is composed of tiny particles called atoms
Atoms of a given element are identical in size, mass, and other properties; atoms of different elements are not the same
Atoms cannot be subdivided, created, or destroyed
Atoms of different elements combine in simple whole-number ratios to form compounds
In reactions, atoms are combined, separated, or rearranged

4. 3.1 The atom Dalton’s Atomic Theory
Idea was supported because his points were backed by three scientific laws
Law of Conservation of Mass
Matter is neither created nor destroyed but conserved
Law of Definite Proportions
Compounds are contain the same elements in exactly the same proportions by mass regardless of the size of sample
Law of Multiple Proportions
Two compounds are made of the same two elements, then the ratio of the masses of the first element when combined with a set mass of the second will always be small whole numbers

5. 3.1 The atom Additional links

6. 3.2 The structure of the Atom
Dalton defined the atom as a small, indivisible particle that makes up all matter
Scientist in the late 1800’s began to test the points of Dalton’s Theory
Found that atoms were composed of smaller particles and that atoms of the same element might not always be identical
Subatomic particles are protons, neutrons, and electrons

7. Discovery of electron JJ Thompson (1897)
Discovery was made through an experiment known as the cathode ray experiment
Changed model to the Plum Pudding Model of the atom

8. Cathode ray experiment

9. Charge of electron
Thompson test the tube with various elements and concluded that all contained electrons
Robert Millikan (1909)
Measured the charge of the electron
Based upon this information, the charge to mass ratio of the electron was found to be about 1/1837 the mass of a hydrogen atom
Since atoms are neutral, there must also be a positive charge
Since the electron’s mass is so small, there must be another particle to make up its mass

10. Discovery of the nucleus
Ernest Rutherford (1911)
Discovery was made by an experiment known as the Gold Foil Experiment
Assisted by Hans Geiger and Ernest Marsden
Changed the model of the atom based upon his findings

11. Gold Foil Experiment
From the experiment, Rutherford was able to conclude that the majority of the mass of the atom is located in a dense, positively charged center
Also concluded that the volume of the nucleus is relatively small compared to the volume of the atom

12. Composition of the nucleus
Nucleus is composed of two subatomic particles
Protons and neutrons
These two subatomic particles have relatively equal masses
Atoms have equal number of protons and electrons
Forces in the Nucleus
Short range force between proton-proton, proton-neutron, and neutron-neutron
Known as Nuclear Force

13. 3.3 Counting Atoms Atomic Number Isotopes
Number that identifies the number of protons
By default, also identifies the number of electrons for neutral atom
Isotopes
Atoms of the same element but have different number of neutrons

14. Counting atoms Mass number Atomic Mass
Sum of the protons and neutrons for a single atom of an element
Atomic Mass
Weighted average of the isotopes for all the atoms of an element

15. Designating isotopes Methods for identifying isotopes Hyphen Notation
Name of element – mass number
Nuclear Symbol
Mass number over atomic number then the symbol of the element
Number of neutrons = mass number – atomic number

16. Calculating atomic mass
Weighted average of the isotopes
Mass number of each isotope multiplied by the percentage that the isotope exists in nature then added together
EX.
Copper-63 (69.15%) and Copper-65 (30.85%)
63 x =
65 x =
= (63.62)

17. Relating mass to number of atoms
Mole
Amount of a substance that contains the same number of particles as Carbon-12
Avogadro’s number
The number of particles in 1 mole which is equal to 6.02 x 1023
Molar Mass
Number of grams equal to 1 mole of a substance
For an element, this is equal to the atomic mass expressed in grams

18. Basic Conversions