1. The Structure of the Atom
Chapter Four
The Structure of the Atom

2. 4.1 Early Theories of Matter
Philosophers believed matter was made of earth, water, air, and fire.

3. Democritus ( BC)
was first to propose that matter was made up of atomos, which could not be further divided
Atoms have different sizes and shapes giving them different properties

4. John Dalton (1766-1844) Dalton’s Atomic Theory-
Referred to the atom as a “hard sphere”

5. All matter is made up of atoms
Atoms of the same element are identical
Atoms cannot be created, divided, or destroyed
Atoms combine in certain reactions to form compounds
In chemical reactions atoms are separated, combined, or rearranged

6. 4.2 Subatomic Particles and the Nuclear Atom
1879-Crookes invented the cathode ray
Cathode Ray- a ray of radiation that originates from the cathode and travels to the anode of a cathode ray tube

7. Led to invention of television

8. By end of the 1800s scientists concluded that cathode rays were a stream of charged particles
Particle carried a negative charge
Electron- negatively charged particle

9. Evolution of the Atom
J.J. Thomson ( ) determined mass-to-charge ratio of the electron
Determined that charged particle mass was less than that of smallest element, hydrogen

10. Meant that atoms were made of smaller particles, disproving Dalton’s theory
Created Plum-pudding model of the atom
Robert Millikan ( ) determined that an electron has a negative charge

11. Thomson’s Plum-pudding or chocolate-chip cookie dough model of the atom
Proposed that negatively charged electrons (chips) were distributed through a “dough” of positive charge.

12. The Nuclear Atom
Ernest Rutherford ( ) used a gold foil experiment to discover existence of nucleus
Nucleus- dense region in center of atom which is positively charged and contains virtually all of its mass

13. Used a gold foil experiment to see if positive alpha particles would be deflected by the electrons in the atom.

14. Since the positive charge was thought to be spread out, he thought it would not alter the path of the alpha particles.

15. Amazingly some were deflected at large angles which meant there must be a concentrated positive area.

16. Completing the Atom- The Discovery of Protons and Neutrons
Rutherford refined concept of nucleus to include protons and neutrons

17. Bohr (1913) Electrons orbit the nucleus
Orbitals have a set size and energy level
Lowest energy level is the smallest orbit

18. Defining the Atom
Atom- the smallest particle of an element that retains the properties of the element

19. How big is an atom?
consider this:
world population in 2000:
# of atoms in a single copper penny:

20. Proton- subatomic particle carrying a positive charge
Electron -subatomic particle carrying a negative charge
Neutron- has a mass nearly equal to a proton, but carries no charge (neutral)

21. Protons, Electrons and Neutrons
+ve charge
Electrons
-ve Charge
Neutrons
0 charge

22. Mass
Protons
Electrons
Neutrons

23. 4.3 How Atoms Differ
Atomic Number- number of protons in an atom
Atomic number= # protons = # electrons

24. Mass of Individual Atoms
Atoms have extremely small masses which are hard to work with, so scientists use a standard for comparison
Standard used is a Carbon-12 atom

25. Carbon-12 atom has mass of 12 atomic mass units
1 atomic mass unit (amu) is nearly equal to mass of 1 proton or 1 neutron
Atomic mass is average mass of isotopes of element

26. Symbolic Notation Mass number
atomic number = number of protons = number of electrons
Symbol
Mass number

27. Isotope identification
Uranium – 238
Element name mass #

28. Mass Number= # protons + # neutrons
Isotope- atoms with same number of protons but different numbers of neutrons

29. Percent Abundance
The chance that it will be found in nature

30. Average Atomic Mass
The average atomic mass is equivalent to the most abundant isotope

32. Isotopes 
Potassium-39
Potassium-40
Potassium-41
Protons 19
Electrons
Neutrons 20 21 22
Symbolic Notation K
40 K
41 K

33. Calculating Atomic Mass
Isotopes of elements exist in nature in varying amounts
Atomic Mass = sum of % abundance x atomic mass for each isotope

34. Calculate the atomic mass unit of chlorine, whose percent abundance is 75% of chlorine-35, and 25% of chlorine -37.
Ex: Chlorine:
Isotopes
Chlorine-35 x (75%) = amu
Chlorine-37 (25%) = 9.25 amu
Atomic Mass = 35.5 amu ( amu)

35. STOP!!!