1. Atoms

2. Let’s Review! Matter is…
Anything that has mass and takes up space
All matter is made of elements – substances that cannot be broken down
And elements are made of – ATOMS!
The atom is the SMALLEST unit of matter

3. ATOMS ARE EVERYWHERE! They are in the air you breathe, the chair you’re sitting on, and the clothes you’re wearing.

4. Section 1
Atomic Theory

5. Democritus (400 B.C.) Ancient Greece
Theorized that matter could not be divided infinitely, you had to reach a smallest piece
Atomos: indivisible or can’t be cut (becomes atom)
Since this was ancient Greece, he had no proof and few people believed it.

6. John Dalton (1808)
People were more accepting of his atomic theory because he had evidence
His theory:
all matter is made of atoms (small solid spheres - Billiard Ball Model)
atoms are indivisible and indestructible
atoms of the same element are exactly alike
atoms of different elements are different
compounds are formed by joining two or more atoms

7. John Dalton (1808)
Dalton’s theory of atoms supported the Law of Definite Proportions
Law of Definite Proportions: A chemical compound is made up of the same percentage of elements.
This means that elements combine is whole number ratios when they react
Water is always made up of 2 parts hydrogen to 1 part oxygen giving us H2O
This is true for all chemical compounds!

8. Sir J. J. Thomson (1897)
Conducted experiments with cathode rays that showed atoms could be divided
His experiments showed negatively charged particles that came from inside atoms
These are electrons!

9. Sir J. J. Thomson (1897)
A cathode ray has two metal plates at the end of a vacuum tube
Cathode, which has a negative charge
Anode, which has a positive charge
When voltage is applied across the plates, a glowing beam comes from the cathode and strikes the anode
Since it was a vacuum tube (a tube with all the air vacuumed out), he could see that the beam came from the negative end of the cathode

10. Sir J. J. Thomson (1897)
Successfully separated negative particles, but could not separate the positive particle
Proposed that electrons are spread throughout the atom like blueberries in a muffin
Called the Plum Pudding Model

11. Ernest Rutherford (1911)
Discovered the nucleus using the gold foil experiments
Fired positively charged particles at a sheet of gold foil
Most went through unaffected, some bounced away
This suggested that an atom’s positive charge was concentrated at the center of the atom
This is the NUCLEUS!

12. Ernest Rutherford
Stated that electrons are scattered around the atom with mostly empty space between them and the nucleus
Compared to the atom, the nucleus is very small, like a marble on a football field!

13. Niels Bohr (1913)
Proposed that electrons are arranged in circular energy levels around the nucleus
Like planets orbiting the sun
When electrons gain energy, they “jump” from a lower level to a higher, a loss of energy causes it to “fall” from a higher level to a lower

14. Erwin Schrodinger (1926) & James Chadwick (1932)
His model does not define the exact path of an electron, but predicts its location
Shows the nucleus surrounded by an electron “cloud”
Chadwick discovered the neutron, which has no electrical charge, and the same mass as the proton

15. Modern Atomic Theory Today we know that atoms have A nucleus with And
Protons and Neutrons
And
Electrons that orbit the nucleus and behave like waves on a string

16. Section 2
Atomic Structure

17. Subatomic Particles Subatomic: lower (or smaller) than an atom
Protons and electrons have an electrical charge
Protons: positive charge
Neutrons: no charge (neutral)
Electrons: negative charge

18. Mass and Volume The nucleus makes up 99.99% of the mass of the atom.
However, the nucleus is 1/100,000 of the volume of an atom.
The volume is determined by the electron cloud.

19. Mass and Volume
Since subatomic particles are so small they cannot be measured in grams
They are measured in atomic mass units or amu
1 amu = 1.61x10-24 g !
1 g is about the mass of a paper clip!

20. protons p+ positive nucleus 1 amu neutrons no neutral electrons e-
name
symbol
charge
location
mass
protons p+
positive
nucleus
1 amu
neutrons no
neutral
electrons e-
negative
electron shell
.0006 amu

21. Electrons Fill energy levels around the nucleus
Each level only holds so many electrons
Valence Electrons: electrons that fill the outermost level

22. Atomic Number Atomic Number: number of protons found in the nucleus
The # of protons in an atom is unique to each element and is how we identify an element – it NEVER CHANGES
In a neutral atom, the number of electrons equals the number of protons – meaning there will be NO overall charge on the atom

23. Atomic Mass Atomic Mass: total mass of all atom’s components
Mass Number: total number of neutrons and protons in an atom
subtract atomic # from atomic mass to find number of neutrons
The number of neutrons in an atom can change, which means the atomic mass can change!

24. Let’s Try! Atomic # Mass Protons Neutrons Electrons 14 28 Atomic #3 7 4
Atomic #
Mass
Protons
Neutrons
Electrons 9 19 10
Atomic #
Mass
Protons
Neutrons
Electrons 90
232
142

25. Isotopes
Isotope: atoms of the same element with a different number of neutrons, and therefore different masses.
Normally 1-2 stable isotopes for an element, and the atomic mass of the most common isotope is listed in your periodic table
All others are unstable (they fall apart) through radioactive decay

26. Representing Isotopes
Isotopes have the SAME chemical and physical properties.
Isotopes can be represented in a number of ways called isotope notation
Element symbol or name with atomic mass
Element symbol with atomic number and atomic mass