1. Atoms
Chapter 4

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. Atomic Theory
Section 1

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 started to accept his idea of atoms because of his experiments
He thought:
all elements are made of atoms (indivisible and indestructible)
atoms of the same element are exactly alike
atoms of different elements are different

7. Law of Definite Proportions
A chemical compound is made up of the same percentage of elements.
For example, water is always made up of 2 parts hydrogen to 1 part oxygen.
Therefore we have the equation
H2O
This is true for all chemical compounds!
This law supported Dalton’s theory of atoms.

8. Sir J. J. Thomson (1897)
Conducted an experiment that showed atoms could be divided
Used cathode ray experiments to show negatively charged particles that came from inside atoms
These are electrons!

9. Cathode Rays There are 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. Plum Pudding
He proposed that electrons are spread throughout the atom just like blueberries are in a muffin
Plum Pudding Model

11. Thomson Found negative particles could come from neutral elements
Atom is made of smaller things (+ & -), and is divisible
Successfully separated negative particles (electrons) but could not separate the positive particle

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

13. Ernest Rutherford
Electrons are scattered near the outside of the atom with mostly empty space between the nucleus and the electrons
Compared to the atom, the nucleus is very small
Rutherford’s experiments led to a new model of the atom

14. Niels Bohr (1913)
Rutherford proposed that electrons orbited the nucleus
Similar to planets orbiting the sun
Bohr discovered the electrons are actually in energy levels (orbitals)
Their path around the nucleus is limited to their energy
Only certain electrons can be in certain levels

15. Modern Atomic Theory Today we know that atoms have A nucleus with
Protons
Neutrons
Electrons that orbit the nucleus

16. TUESDAY, March 8th AFTER THE QUIZ: You need: Guided Reading Notes
Calculator
Periodic Table

17. Atomic Structure
Section 2

18. Subatomic Particles Subatomic: lower (or smaller) than an atom
Protons: positive charge
Neutrons: no charge (neutral)
Electrons: negative charge
Protons and neutrons are both in the nucleus of the atom.
Electrons are buzzing around the outside in the electron cloud

19. 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.

20. Mass of Particles
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!

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

22. How do we know how many protons, neutrons, and electrons are in an atom?

23. Atomic Number & Atomic Mass
number of protons found in the nucleus (therefore, it is also the number of electrons orbiting the nucleus)
Atomic Mass:
weight of the protons and the neutrons combined.

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. Charges
The # of protons in an atom is unique to each element (this is how they are identified)
If an atom is stable, there will be an equal number of protons and electrons
This means there will be NO overall charge on the atom
The number of neutrons in an atom can vary also!

26. Electrons Electrons fill orbitals around the nucleus
Each orbital can only hold so many electrons
The number of levels filled depends on how many electrons the atom has
The electrons that fill the outermost orbital are called valence electrons

27. Isotopes
The atomic number, or number of protons, will never change, but the number of neutrons CAN change
Isotope: Any two atoms of the same element with a different number of neutrons.
Isotopes have the SAME chemical and physical properties.
There are normally one or two stable isotopes for an element
All others are unstable (they fall apart)
radioactive decay

28. Representing Isotopes
Isotopes are shown in isotope notation
Since isotopes have a different atomic mass, you must write the symbol of an element AND write its atomic mass

30. MONDAY, March 14th
PICK a PARTNER that you are SURE you will succeed with in lab today.
You will walk around the classroom together and compile information about atoms for your data table.
THEN return to your seat and FINISH the data table (Atomic #, MASS, ELEMENT, SYMBOL, ISOTOPE?)
ANSWER Analysis Questions.
SUBMIT BEFORE end of PERIOD.

31. RED- PROTONS
BLUE-NEUTRONS
YELLOW- ELECTRONS