1. Building blocks of matter
Atoms
Building blocks of matter

2. Foundations of atomic theory
Laws
Law of conservation of mass
Mass is neither created nor destroyed
What goes in must come out
Law of definite proportions
A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound
Salt is salt and is always salt

3. One more law Law of multiple proportion
If two or more different compounds are composed of the same two elements, than the ratio of the masses of the second element combined with a certain mass of the first element is always a ration of small whole numbers.

4. Dalton’s atomic theory
Dalton was an English schoolteacher and in 1808 he proposed an idea to explain all these laws. An idea known as Dalton’s atomic theory.
1. all matter is composed of extremely small particles called atoms.
2. atoms of a given element are identical in size, mass and other properties; atoms of different elements differ in size, mass and other properties.

5. Dalton continued 3. atoms cannot be subdivided, created, or destroyed.
4. atoms of different elements combine in simple whole-number ratios to form chemical compounds.
5. in chemical reactions, atoms are combined, separated, or rearranged.

6. Modern atomic theory
Dalton’s ideas were revolutionary but not all were correct
We know now that atoms can be separated into smaller parts.
We know now that a given element can have atoms with different masses
But we tweaked it and keep some of his ideas
All matter is composed of atoms
Atoms of any one element differ in properties from atoms of another element remain unchanged
end 3-1

7. The structure of the atom
An atom is the smallest particle of an element that retains the chemical properties of that element.
We know the atom has a nucleus found in the middle of the atom made of
Protons
Neutrons
With electrons orbiting on the outside

8. Discovery of the electron
The electron was discovered in the 1800’s by passing electricity through a gas in a sealed tube
This tube is now called a cathode-ray
Cathode is the negative side of a battery
Anode is the positive side

9. Cathode rays and electrons
When scientists started to experiment with cathode rays they noticed that when electricity passed through the tube the tube glowed directly opposite of the cathode.
They decided that these had to be particles and continue to do experiments and observed the following things happened.

10. Observations of the cathode tube
1. an object placed between the cathode and the opposite end of the tube cast a shadow on the glass
2. a paddle wheel placed on rails between the electrodes rolled along the rails form the cathode toward the anode.
These things told scientists that the ray had some mass.

11. Continued
3. cathode rays were deflected by a magnetic field in the same manner as a wire carrying electric current, which was known to have a negative charge
4. the rays were deflected away from a negatively charged object.

12. Joseph John Thomson
Thomson did experiment that revealed the charge of the electron.
His experiments showed that electrons has a great charge for their mass
Millikan found the mass of electrons
Smaller than the smallest hydrogen atom
9.109 x 10 −31

13. Inferences about atoms from experiments
Sense Thomson proved electrons have a negative charge and Millikan showed they don’t have much mass we can infer
1. because atoms are electrically neutral, they must contain a positive charge to balance the negative electrons
2. because electrons have so much less mass than atoms, atoms must contain other particles that account for most of their mass

14. Discovery of the atomic nucleus
Rutherford and Geiger did an experiment.
What they did was shoot alpha particles at a sheet of gold foil.
What they found out shocked them
Some of the alpha particles were deflected right back at the source.
What this told them was that the center of the atom was something solid.

16. More on Rutherford
By using these results Rutherford discovered that the volume of the atom was mostly empty space.
That lead to the realization that there was a tightly pack nucleus.
He also inferred that the nucleus must have a positive charge to balance out the atom

17. What’s in the nucleus
Except for hydrogen all elements have two particles in their nuclei
Protons
Positive charge
Mass of 1.673x 10 −27
About 1836 times greater than that of an electron
Neutrons
For them no charge
Mass of 1.675x 10 −27
Just slightly heavier than a proton

18. Forces in the nucleus As we know like charges repel each other
Therefore we would expect a nucleus with more than one positive proton to be unstable and try to get away from each other.
In the nucleus this is not true.
The closer the protons are together the greater the attraction.
Same holds true for neutrons
And when protons are close to neutrons
These short-range proton-neutron, proton-proton and neutron-neutron forces are referred to as nuclear forces. END 3-2

19. Counting atoms Atomic number (Z)
The number of protons in the nucleus of each atom of that element
The periodic table is step up using these (more on this later)
Identifies the element
Atomic number 6 is Carbon
Carbon will always have an atomic number of 6

20. Isotopes
An isotope is an atom that has a different number of neutrons and therefore has a different atomic mass.
Hydrogen normally has only one proton
Protium= one proton
Deuterium= one proton and one neutron
Tritium= one proton and two neutron
Doc oct uses this to make the mini sun.

21. Mass number
The mass number of an element is very simple to figure out.
It is the sum of all the particles in the nucleus
So its protons plus neutrons.
Changing the mass number does not change the atomic number.
It does not change the identity of the element

22. Designating isotopes
Usually isotopes don’t have distinct names like those of hydrogen.
There are two methods for identifying isotopes
First method is the hyphen notation
Hydrogen-3
Uranium-235
Used in nuclear reactors
Second method is the nuclear symbol
Uranium-235 is written like 92 𝑈
Nuclide is the general term for any isotope of any element
235

23. Relative atomic masses
Used to set a standard of measurements
Uses carbon-12
Carbon-12 is exactly 12 amu
Amu= atomic mass units
Although isotopes have different masses, they do not differ significantly in their chemical behavior.

24. Average atomic masses of elements
The weighted average of the atomic masses of the naturally occurring isotopes of an element.
For example you have a box full of 100 marbles
25 are small with a mass of 2.00 g
75 are large with a mass of 3.00g
Total mass is 25*2+75*3= 275
Then divide by 100 and you get 2.75

25. example Copper has two isotopes Copper-63 which makes up 69.17%
Atomic mass amu
Copper-65 which makes up 30.98%
Atomic mass amu
Now take the decimal values for the percent multiply by the atomic mass and add the results together.
.6917* *
63.55amu

26. Mass to numbers of atoms
The mole
A mole (mol) is the amount of substance that contains as many particles as there are atoms in exactly 12g of carbon-12.
Its like saying a dozen

27. Avogadro’s Number 6.0221367x 10 23 rounded to 6.022x 10 23
Ok but what does this number mean
Well lets take 5 billion people about everyone on earth that knows how to count. Have the start counting atoms at a rate of one atom/second
It would take only about 4 million years to count all the atoms in one mole

28. Molar mass The mass of one mole of a pure substance
Lucky for us for the element the molar mass is the same as the atomic mass
Li=6.94 C= O=16.00
When determining the molar mass of compounds and molecules just add all of them up.
𝐻 2 O= two hydrogen and one oxygen
Molar mass of water = g/mol

29. Gram/Mole conversions
Use molar mass=g/mol
So how many grams is 2.00 mol of He
Look for the molar mass from the periodic table
4.00 grams/mol
2.00 mols *4.00 grams/mols
8.00 grams of He