1. Ch. 5 Atomic Structure and The Periodic Table

2. I. Atomic Model Theories A. Dalton’s Theory (1807) 1. He theorized that an atom was indivisible, uniformly dense sphere. 2. He theorized that all atoms of the same element have the same mass and the same chemical behaviors.

3. I. Atomic Model Theories A. Dalton’s Theory (1807) 3. He theorized that atoms of different elements have different chemical behaviors. 4. He theorized that atoms of different elements combine to form compounds. (Example — H 2 O)

4. I. Atomic Model Theories B. Just How Small Is an Atom? 1. The atom is the smallest particle of an element that retains the properties of the element. 2. It can not be seen by the naked eye.

5. II. Structure of the Nuclear Atom A. Electrons 1. J.J. Thompson devised an experiment to determine the nature of the cathode ray. 2. He built a cathode ray tube and placed a fluorescent screen at the end that would glow when struck by charged particles. 3. When the beam was normal, the center of the screen would glow but, when altered with a magnet or charged plates, the glow would move in the direction that the beam was altered.

6. II. Structure of the Nuclear Atom A. Electrons 4. He concluded that the direction of the beam was determined by the charge of the plate. 5. He concluded that the particles in the cathode ray are subatomic particles that are found in all atoms. 6. He is credited with the discovery of the negatively charged particles called electrons.

7. II. Structure of the Nuclear Atom A. Electrons 7. He theorized that the atom is a dense sphere with a positive charge and also contains negative charged particles. 8. Robert A. Millikan obtained the first accurate measurement of an electron charge. 9. Using a brass atomizer, he sprayed oil drops into an apparatus with charged plates inside (electrons were transferred to oil drops).

8. II. Structure of the Nuclear Atom A. Electrons 10. The charge plates were set to offset the force of gravity. 11. When the forces were equal, the drops were stationary (the drops did not move). 12. An electron carries exactly one unit of negative charge, and its mass 1/1840 the mass of a hydrogen atom.

9. II. Structure of the Nuclear Atom B. Protons and Neutrons 1. Evidence for a positively charged particle was found by E. Goldstein. 2. He observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays. 3. He called them canal rays and concluded that they were composed of positively charged particles.

10. II. Structure of the Nuclear Atom B. Protons and Neutrons 4. We call these particles protons. 5. The first evidence of the third particle was discovered by Walter Bothe and later James Chadwick repeated Bothe’s work. 6. They found a high energy particles with no charge and the same mass as the proton. 7. These particles are called neutrons.

11. II. Structure of the Nuclear Atom C. The Atomic Nucleus 1. Ernst Rutherford designed an experiment to test Thompson’s Model. 2. He used alpha particles to bombard targets made of thin gold foil. 3. A fluorescent screen was placed around the gold foil to detect the particles after they struck the target.

12. II. Structure of the Nuclear Atom C. The Atomic Nucleus 4. He expected that the particles should uniformly pass through the foil undisturbed and some of the particles did pass through but other were deflected.

13. II. Structure of the Nuclear Atom C. The Atomic Nucleus 5. He realized the explanations that accounts for the deflection. a. The atom has a very dense center of positive charge called the nucleus. b. The nucleus contains the protons for the atom and make up more than 99.9% of its mass. c. The electrons move around the nucleus.

14. II. Structure of the Nuclear Atom C. The Atomic Nucleus

15. III. Distinguishing Between Atoms A. Atomic Number 1. The atomic number is the number of protons in an atom. 2. Because the atom is electrically neutral, the number of electrons is equal to the number of protons. 3. The number of protons determines the identity of the element. 4. To solve for the number of neutrons, subtract the atomic mass by the atomic number.

16. III. Distinguishing Between Atoms A. Atomic Number

17. III. Distinguishing Between Atoms B. Mass Number 1. The proton and neutron are equal in mass. 2. The mass of the electron is extremely small, so most of the mass is in the nucleus. 3. It is possible to discuss the mass of one atom, however chemist use the masses of a large groups of atoms.

18. III. Distinguishing Between Atoms B. Mass Number 4. The way to represent this is with the symbol of the element and the mass as a superscript and the atomic number as a subscript. ( Oxygen 16 8 O) 5. It can also be written as oxygen - 16.

19. III. Distinguishing Between Atoms C. Isotopes 1. An element that has the same atomic number but a different mass is called an isotope. 2. Isotopes are chemically alike because they have identical numbers of protons and electrons which are responsible for reactivity. 3. Example Carbon-12 and Carbon-14

20. III. Distinguishing Between Atoms D. Atomic Mass 1. There are two ways of determining masses for atoms of other elements a. Reacting the standard element with the element to be determined b. Mass Spectrometer. 2. Using a mass spectrometer, we can determine the relative amounts and masses of the nuclides for all isotopes of an element.

21. III. Distinguishing Between Atoms D. Atomic Mass 3. Chemist measure the mass of one atom in atomic mass units (AMU). 4. The mass of the atomic particles. a. electron = 9.109 x 10 -28 g = 0.000549 AMU b. proton = 1.672 x 10 -24 g = 1.00 AMU c. neutron = 1.674 x 10 -24 g = 1.00 AMU 5. This is called the atomic mass and is the average of the amounts and masses of all the isotopes of the element.

22. IV. The Periodic Table: Organizing the Elements A. Development of the Periodic Table 1. Dmitri Mendeleev proposed an arrangement. 2. He suggested that the properties of the elements was a function of their atomic masses, however, he believed that similar properties occur after periods that could vary in length.

23. IV. The Periodic Table: Organizing the Elements A. Development of the Periodic Table 3. He also left some blank spots in the order to suggest that there were elements yet to be discovered. 4. He stated that the properties of the elements are a periodic function of their atomic mass and this was called the periodic law. 5. Henry Moseley found exceptions to Mendeleev’s periodic law.

24. IV. The Periodic Table: Organizing the Elements A. Development of the Periodic Table 6. He performed X–ray experiments showed that the nucleus of each element has an integral positive charge which is the atomic number. 7. As a result of Moseley’s work, periodic law was revised and is now based on the atomic number instead of atomic mass. 8. The modern periodic law now states that the properties of the elements are a periodic function of their atomic number.

25. IV. The Periodic Table: Organizing the Elements B. The Modern Periodic Table 1. Certain electron arrangements are repeated periodically as atoms increase in atomic number. 2. All elements in a horizontal row are referred to as a period. 3. All elements in the same vertical column are referred to as a group.

26. IV. The Periodic Table: Organizing the Elements B. The Modern Periodic Table 4. Metals are substances that are hard, shiny, and conduct heat and electricity. 5. Nonmetals are substances that are brittle, dull, solids or gases at room temperature. 6. Nonmetal do not conduct electricity or heat and are insulator.

27. IV. The Periodic Table: Organizing the Elements B. The Modern Periodic Table 7. The following are the metal and nonmetal families. a. Alkali metals b. Alkaline earth metals c. Transition metals d. Lanthanoids e. Actinoids f. Chalcogens g. Halogens h. Noble gases