1. Atoms: The Building Blocks of Matter
Chapter 3
tPHzzYuWy6fYEaX9mQQ8oGr

2. What is an atom?
The smallest particle of an element that retains the chemical properties of that element.
Regions:
Nucleus: very small region located at the center of atom
Outside the nucleus: electrons most in the “electron cloud”
Subatomic Particles:
Neutron, neutral found in nucleus (mass number subtracted by number of protons indicates neutrons)
Proton, positive found in nucleus (atomic number indicates number of protons in the atom)
Electron, negative found surrounding the nucleus (equal number of protons and electrons)

3. John Dalton Dalton’s atomic theory explains all three of these laws.
All elements were composed of atoms and that only whole numbers of atoms can combine to form compounds. The following statements sum up his theory:
All matter is composed of extremely small particles called atoms.
Atoms of an element are identical in size, mass, and other properties.
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.

4. Modern Atomic Theory
Several changes have been made to Dalton’s theory…
Dalton said…Atoms of an element are identical in size, mass, and other properties.
Modern Theory states: Atoms of an element have the same number of protons (atomic number) but can have different masses (because of the different number of neutrons)
Dalton said…Atoms cannot be subdivided, created, or destroyed.
Modern Theory states: Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!

5. Discovery of Subatomic Particles
Late 1800s, JJ Thomson, used an cathode ray tube to determine the existence of negative electrons by passing an electric current passed through low pressure gases.
Cathode-ray tube experiment
Cathode (- charged metal disk side of the tube)
When electricity was passed through the tube, the directly opposite side of the cathode would glow.
The assumed the glow was due to the particle stream and called this glow a “cathode ray.”
Particles that compose cathode rays are negatively charged (since they were attracted to the + anode end)
Ultimately, this is how an electron was found!

7. The Nuclei Ernest Rutherford’s Gold Foil Experiment
Fired helium nuclei (alpha particles) at a thin sheet of gold foil
A screen record where they hit
Only about 1 in 8000 of the alpha particles bounced off the gold foil. Not what they expected.
Concluded that “some powerful source within the atom deflected the alpha particles” and the force must be caused by a very densely packed bundle of matter with a positive charge. This positively charged bundle was then called the nucleus. He found the nucleus to be a very small fraction of the overall size of the atom.
Football field atom – marble nucleus
Niels Bohr, proposed a model that resembled the sun and the planets.

8. Rutherford’s Findings
Most of the particles passed right through
A few particles were deflected
Only about 1 in 8000 alpha particles bounced off the gold foil.
VERY FEW were greatly deflected
Conclusion:
The nucleus is small
The nucleus is dense
The nucleus is positively charged

9. The Nucleus Protons Neutrons
Positive charge equal in magnitude to the negative charge of an electron.
1.673 x kg
Neutrons
Electrically neutral charge.
1.675 x kg
Neutrons are slightly larger in mass than protons
The nuclei of atoms of different elements differ in their number of protons, thus the number of protons determines the atom’s identity.

11. Forces
Generally, particles that have the same charge repel one another.
More than one proton in the nucleus you would expect to be unstable. However, it’s the opposite. When two or more protons are extremely close to each other, they have a strong attraction.
Nuclear force is known as short-range proton- neutron, proton-proton, and neutron-neutron forces hold the nuclear particles together.

12. Bohr
Developed the planetary model and proposed that the electrons orbit around the nucleus in a circular part at a constant speed.
Incorrect!!

13. Schrodinger Quantum Mechanical Model The model we use today!!
Also called the electron cloud model.
The electrons don’t have a distinctive orbit, they have a probability region where they can be found.
Think about a fly buzzing around the room…very similar to how an electron would act!

14. Subatomic Particles Charge Mass Location Proton + 1 amu Nucleus
Neutron
None
Electron -
Almost nothing…
Outside Nucleus
(in motion)

15. Determining Subatomic Particles
For any element: # of Protons = Atomic # # of Electrons = # of Protons = Atomic # # of Neutrons = Mass # - Atomic #

16. Work on Part 1 Worksheet

17. Atomic Number
Atoms of different elements have different number of protons.
The atomic number of an element is the number of protons of each atom of that element.
Hydrogen, H, has atomic number 1.
One proton in each atom of Hydrogen.

18. Isotopes Atoms of the same element that have different masses.
Remember, protons and neutrons make up the majority of the mass of an atom.
All isotopes of hydrogen have the same number of protons and electrons but vary in neutrons.
Most elements consist of isotope mixtures.
Tin has 10 stable isotopes, the most of any element.

19. Isotopes
Just because every atom of Hydrogen has one proton…it doesn’t mean they have the same number of neutrons.
Actually, three types of Hydrogen atoms exist!
Protium
Accounts for % of hydrogen atoms on Earth
No neutrons, 1 proton
Total mass of 1
Deuterium
Accounts for % of hydrogen atoms on Earth
1 neutron, 1 proton
Total mass of 2
Tritium
Radioactive and is rarely found on Earth only in trace amounts.
2 neutrons, 1 proton
Total mass of 3

20. Identifying Isotopes Two Methods:
Hyphen-Notation – where the mass number appears after the element with a hyphen.
Nuclear-Symbol – the composition that shows the mass number and atomic number
Nuclide – is a general term for a specific isotope of an element

21. Isotope Worksheet & Homework

22. Review Last Class https://youtu. be/FSyAehMdpyI

23. Law of Conservation of Mass
Mass is neither created nor destroyed during ordinary chemical reactions or physical changes.

24. 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.

25. Law of Multiple Proportions
If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers.

26. Atomic Mass
Scientists use the relative atomic masses because they are most convenient.
Standard for this measure is the carbon-12.

27. Calculating Average Atomic Mass
The weighted average of the atomic masses of the naturally occurring isotopes of an element.
Depends on mass and relative abundance
Calculated by multiplying the atomic mass of each isotope by its relative abundance and adding the results.

28. Average Atomic Mass
Suppose you have a box containing two sizes of marbles. If 25% of the marbles have masses of 2.00 g each and 75% have masses of 3.00 g each, how is the weighted average calculated? Suppose you have 100 marbles.
25% of 100 = 25 marbles
75% of 100 = 75 marbles
25 marbles x 2.00 g = 50 g
75 marbles x 3.00 g = 225 g
50 g g = 275 g ÷ 100 marbles = 2.75 g per marble

29. Work on Worksheet

30. Review What chemical laws can be explained by Dalton’s theory?
Three compounds containing potassium and oxygen are compared. Analysis shows that for each g of O, the compounds have g g, and g of K, respectively. Show how these data support the law of multiple proportions.