1. Chapter 3 Atoms: The Building Blocks of Matter

2. Atomic Theory of Matter
The theory that atoms are the building blocks of matter
by John Dalton.
Figure 2.1 John Dalton ( )

3. Law of Conservation of Mass
The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.

4. Dalton’s Postulates
Each element is composed of extremely small particles called atoms.

5. Dalton’s Postulates
All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.

6. Dalton’s Postulates
Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

7. Dalton’s Postulates
Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

8. The Electron
Figure 2.4
Streams of negatively charged particles were found from cathode tubes.
J. J. Thompson is credited with their discovery (1897).

9. The Atom, circa 1900: “Plum pudding” model, put forward by Thompson.
Positive sphere of matter with negative electrons imbedded in it.
Figure 2.9

10. Discovery of the Nucleus
Ernest Rutherford shot particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.
Figure 2.10

11. Other Subatomic Particles
Protons were discovered by Rutherford in 1919.
Neutrons were discovered by James Chadwick in 1932.

12. Subatomic Particles
Protons and electrons are the only particles that have a charge.
Protons and neutrons have essentially the same mass.
The mass of an electron is so small we ignore it.
Table 2.1

13. Symbols of Elements
Elements are symbolized by one or two letters.

14. Atomic Number
Number of protons = The atomic number

15. Atomic Mass
The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

16. Isotopes:
Atoms of the same element with different masses.
Isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C

17. Atomic Mass
Atomic and molecular masses can be measured with great accuracy with a mass spectrometer.
Figure 2.13

18. Avogadro’s Number
6.02 x 1023 – is the number of particles in exactly one mole of a pure substance.
Conversion factor!
6.02 x 1023 particles = 1 mole

19. If I have 3. 45 moles of hydrogen, how many particles do I have. 6
If I have 3.45 moles of hydrogen, how many particles do I have? x = 1 Mole

20. Molar Mass
The molar mass of an element is the mass number for the element that we find on the periodic table
Look on the periodic chart and find the mass of the element – that equals 1 mole.

21. What is the mass of NaOH?

22. What is the mass of (NH4)2SO4

23. Moles The mole is the SI unit for amount of substance
A mole (abbreviated mol) is the amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12.

24. Given 3.5 mol of Cu – what’s the mass?
Atomic mass = 1 mole
Atomic mass of copper = grams
grams Cu = 1 mole Cu
Conversions!

25. Given 2.3 mol of He, what is the mass?

26. Given 11.9 g of Al, what is the moles?
From the periodic chart:
grams of Al = 1 mole
Now convert!

27. Given 23.5 g of O, how many moles?

28. Conversions with Avogadro’s Number
2 Step conversions
You can now convert from Moles to grams and then to number of atoms.
Let work some examples!

29. Given 3.2 g of Li, how many atoms do I have?
From the periodic chart:
6.941 grams of Li = 1 Mole
6.022 x 1023 atoms = 1 Mole

30. How many grams of silver are in 3.01 x 1023 atoms of silver?
From the periodic chart:
grams Ag = 1 Mole
6.022 x 1023 atoms = 1 Mole