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

2. The Atom: From Philosophical Idea to Scientific Theory
Section 3-1
The Atom: From Philosophical Idea to Scientific Theory

3. Chemical Reaction
A reaction is a change in which one or more substances are converted into different substances.
The reactants are the starting materials.
The products are the ending materials.

4. The Law of the conservation of mass
Matter can neither be created nor destroyed in ordinary chemical reactions.
This means that mass of reactants equals mass of products in a closed system.

5. Law of Conservation of Energy
Energy cannot be created or destroyed.
It can be changed from one form to another.

6. Law of Definite Proportions
A chemical compound contains the same elements in exactly the same proportions by mass regardless of how large a sample or source of the compound.

7. Example: NaCl 100. g sample 39.3% is sodium 60.7% is chlorine

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

9. Example: CO and CO2
In CO 1.00 g of C 1.33 g of O Ratio of oxygens 1.33/2.66 = 1/2
In CO g of C 2.66 g of O

10. Concept of Atom 1. Democritus Greek philosopher (400 B.C.)
called atomos which means “indivisible”.
Aristotle thought matter was continuous. Atoms were ignored for 2000 years.

11. Concept of Atom 2. John Dalton English School Teacher (1800’s)
Proposed explanation of the laws
that existed at that time.
Called Dalton’s Atomic Theory.

12. a) All matter is made of atoms.
Dalton's Atomic Theory
 a) All matter is made of atoms.
b) All atoms of a given element are identical in size, mass and properties.
c) Atoms cannot be subdivided, created or destroyed.

13. Dalton's Atomic Theory
d)Atoms of different elements combine in simple whole number ratios to form chemical compounds.
e) In chemical reactions, atoms are combined, separated or rearranged.

14. Dalton's Atomic Theory
Dalton turned Democritus’ idea into a theory that can be tested.
Testing however showed that his theory needed to be modified.

15. The Structure of the Atom
Section 3.2
The Structure of the Atom

16. Atom
The smallest unit of an element that maintains the properties of that element

17. Parts of the Atom 1)The nucleus
Small region that is dense. Contains protons and neutrons
2)The electron cloud
The region surrounding the nucleus of an atom where electrons are likely to be found.

18. Electron Negatively charge particles found in the electron cloud
Mass: x kg
Electric Charge: x coulombs

19. Proton Positive charge particles found in the nucleus
Mass x 10−27 kg (or about times the mass of an electron)
Electric charge of +1.6 x 10−19 coulomb

20. Neutron Neutral particles found in the nucleus
Mass: x kg (slightly more than a proton)
Electric charge: no net electric charge

21. Thomson's Cathode–Ray Tube Experiment
Time frame was 1897
The discovery was the electron
The scientist was J. J. Thomson

22. Thomson's Cathode–Ray Tube Experiment

23. Thomson's Cathode–Ray Tube Experiment
Concluded that the negatively charged particles in the cathode–ray tube were fundamental particles which are present in all matter.
Called the tiny particles electrons.

24. J.J. Thomson’s Atomic Theory
3. Plum Pudding Theory
Atoms of elements are divisible.
Electrons are small particles that can leave the atom.
Electrons are spread evenly through a positive charge in the rest of the atom like plums in pudding.

25. Millikan's Oil Drop Experiment
In 1909, American physicist Robert Millikan determined the charge of the electron with his famous oil drop experiment.

26. Millikan's Oil Drop Experiment

27. Rutherford's Gold Foil Experiment
Time frame: 1911
New Zealand
Discovery: Nucleus is dense and does not take up much space

28. Rutherford's Gold Foil Experiment
Problem: Mass and charge were believed to be uniformly distributed throughout an atom.
Rutherford's team expected positively charged alpha particles to pass through a sheet of thin gold foil with only a slight deflection.

29. Rutherford's Gold Foil Experiment

30. Rutherford's Gold Foil Experiment
One in 8,000 particles, however, bounced back toward the source. Why?

31. Rutherford's Gold Foil Experiment

32. Rutherford's Gold Foil Experiment
Conclusion: The rebounded particles must have encountered a very densely packed bundle of matter with a positive electric charge.
After much thought, decided nucleus existed.

33. Niels Bohr Atomic Theory
Student of Rutherford
4. Called Planetary theory
Electrons travel around the nucleus much like planets travel around the sun.
Nucleus is very small, dense, positively charged object.

34. Nuclear Forces
They are the interaction that binds protons and neutrons, protons and protons, and neutrons and neutrons together in a nucleus.

35. Nuclear Forces
In stable nuclei, the nuclear force is stronger than the electric repulsion between the protons alone and holds the nucleus together.
In unstable nuclei, the number of protons and neutrons is not balanced, causing the nuclear force to weaken and the nucleus to decay.

36. Forces in Atoms
The forces acting between atomic and subatomic particles.
Gravitational
Electromagnetic
Strong nuclear
Weak nuclear

37. Size of Atoms
Radius of an atom is the distance from center to the outer portion of the electron cloud.
Use picometer (pm = 1 x )
Atomic radii range from 40 to 270 pm
If nucleus were marble, atom would be a football field.

38. Section Three
Counting Atoms

39. Atomic Number Z - The number of protons in the nucleus of an atom.
The atomic number is the same for all atoms of an element
Atomic No. = No. of protons = No. of electrons

40. Mass Number
A - The sum of the numbers of protons and neutrons in the nucleus of an atom.
Because electrons have very little mass, the mass number is very close to the average atomic mass of the element.
Mass No. = no. of neutrons + no. of protons

41. Isotopes and Nuclides
Isotopes are atoms of the same element that have different masses.
Nuclide is the general term for any isotope of any element, based on the number of protons and neutrons in its nucleus.

42. Hydrogen Isotopes Atomic No. No. of Mass No. (No. of Protons) Neutrons
Protium
(Hydrogen-1)
Deuterium
(Hydrogen-2)
Tritium
(Hydrogen-3)

43. Nuclides (Isotopes)
Nuclear No. of No. of No. of Symbol Protons Electrons Neutrons
Helium-3
Helium-4
Sodium-23
Carbon-12

44. Nuclides (Isotopes)

45. Relative Atomic Mass Mass of an atom compared to a standard.
Standard is atomic mass of carbon-12 atom (12 amu)
1 atomic mass unit = 1/12 the mass of a carbon-12 atom (1 amu)

46. Average Atomic Mass
The average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element.
Atomic mass on periodic table is average atomic mass.

47. Average Atomic Mass
Example: 25 marbles weigh 2.00 g each 75 marbles weigh 3.00 g each 25 x 2.00g = 50.0g 75 x 3.00g = 225 g 100 marbles =275 g Average weight for each marble is 2.75 g

48. The Mole
A mole is the amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon–12.
A mole is a quantity like a dozen.

49. Avogadro's Number
The number of particles in 1 mole equals x 1023 per mole

50. Molar Mass Molar mass is the mass in grams of one mole of a substance
A mole of any other kind of atom or molecule will always contain the same number of particles.

51. Comparison of Terms 1 mole of Iron atoms = 55.85 g
1 mole of Iron has x 1023 atoms
1 atom of Iron = amu
1 atom of Iron-55 = 55 amu
1 mole of Iron-55 = 55 g

52. (of terms that is. We still have the math problems to go.)
The End
(of terms that is. We still have the math problems to go.)