Chapter 4 The Structure of the Atom

Transcription

1 Chapter 4 The Structure of the Atom Read pg Early Theories of Matter The Philosophers Democritus Artistotle - Artistotle s influence so great and the science so primitive (lacking!) his denial for the existence of atoms went largely unchallenged for years. John Dalton - Revived and revised ideas bases on research Studied numerous reactions.. Atoms explain the Law of Conservation of mass! 1

2 Dalton has three errors in his theory: Defining the Atom 4.2 Subatomic Particles and The Nuclear Atom Fig. 4-7 The Cathode Ray Tube They determined that the ray s negative particles were found in all forms of matter because Fig. 4-8: Multiple Experiments helped determine the properties of cathode rays. A. B. C. 2

3 JJ Thomson ( ) 1890 s completed a series of experiments that determined. - Compared - SHOCKING! Because Robert Milikan ( ) 1909 Some questions Thomson s and Milikan s work raised JJ Thomson s Model: Fig. 4-8 The Nuclear Atom Ernest Rutherford ( ) Notes about the experimental set-up: 3

4 Fig Expected Fig Set-up Fig Explanation Results & results of results Rutherford s Model: Completing the Atom The Discovery of Protons and Neutrons Atom Basics: Table 4-1: Properties of Subatomic Particles Particle Symbol Location Relative electrical charge Actual Charge (C) (in class) Relative mass Actual mass (g) Electron Proton Neutron Read pg How Atoms Differ Henry Mosely ( ) Atomic number 4

5 Ex. 4-1 Composition of Several Elements Element Atomic Number Protons Electrons a. Pb 82 b. 8 c. 30 Isotopes and Mass Number Isotopes - Chemical behavior is determined by the number of. Mass number Fig Nuclear Symbol Notation Ag Ag Fig Three naturally occurring isotopes of potassium, K. Protons Neutrons Electons Potassium-39 Potassium-40 Potassium-41 Ex. 4-2 Determine the number of protons, neutrons, and electrons in the isotope of neon. Name the isotope and give its symbol (IN CLASS EXAMPLES:) Nuclear Symbol Isotope # Proton # Neutrons # Electrons Mass # Charge Na 1? H 16 8 O 18? O Na

6 Mass of Individual Atoms Atomic mass unit Table 4-2 Masses of Subatomic Particles Particle Electron Proton Neutron Mass (amu) Atomic mass Ex. 4-3 Calculating Atomic mass Given the data in the table at the right, calculate the atomic mass of unknown element X. Then, identify the unknown element, which is used medically to treat some mental disorders. Isotope Mass (amu) Percent Abundance 6 X % 7 X % Other Element examples: Fluorine ( amu) Bromine ( amu) (IN CLASS EXAMPLES:) Isotope Mass of Isotope Abundance C % C % C Trace Isotope Mass of Isotope Abundance Cu % Cu-65?? Unstable Nuclei and Radioactive Decay Radioactivity (ordinary) Chemical reaction - 6

7 Nuclear reaction - Radiation - Nuclei are radioactive because Radioactive decay - Types of Radiation Alpha Radiation Ra + He 88 2 Beta Radiation 14 0 C Gamma Radiation U + He In Class Practice: Kr Rb + He + Pu Pb He Am Np How about Kr undergoing an alpha and two beta decays? 36 Fig Radiation and Charge complete the diagram. 7

8 Nuclear Stability Table 4-3 Characteristics of Alpha, Beta, and Gamma Radiation Radiation Type Alpha Beta Gamma Composition (in class) Symbol Mass (amu) Charge Ch. 5 Electrons in Atoms Read. Pg (be sure to read objectives on pg. 116) The Nuclear Atom and Unanswered Questions Rutherford s model lacked details about In the early 1900 s Wave Nature of Light Electromagnetic radiation Ex s Wavelength - Frequency Amplitude c = 8

9 Fig. 5-3 Wavelength vs. Frequency Continuous spectrum Fig. 5-5 Electromagnetic spectrum Ex. 5-1 Calculating Wavelength of an EM Wave Microwaves are used to transmit information. What is the wavelength of a microwave having a frequency of 2.44 x 10 9 Hz? Particle Nature of Light - The wave model of light cannot explain The quantum concept Max Plank Quantum Prior experience had lead scientists to believe that energy could be absorbed and emitted in continual varying quantities, with no minimum limit to the amount. Actually, The photoelectric effect (You do not need to understand this! Just know what a photon is.) E = Photon E photon = KEY: Einstein was able to explain the photoelectric effect by 9

10 Atomic Emission Spectra - Each element s atomic emission spectrum is - The fact that only certain colors appear in an elements atomic emission spectrum means that 5.2 Quantum Theory and the Atom Bohr Model of the Atom -answered the question why - correctly predicted Energy states of hydrogen - Proposed that Ground state Excited state Ionized state (in class) - - Suggested that An explanation of hydrogen s line spectrum n = 1 ΔE = n = 5 n = 4 n = 3 n = 2 n = 1 Bohr s Model didn t account for 10

11 Read pg The Quantum Mechanical Model of the Atom Electrons as Waves Louis De Broglie ( ) = Key: electrons have properties. The Heisenberg Uncertainty Principle Heisenberg uncertainty Principle The Schrodinger wave equation Quantum mechanical model of the atom Atomic orbital In Class Summary (cont. to next page for notes) Wave Nature Energy Particle Nature of Energy Bohr Heisenberg Schrodinger - 11

12 Hydrogen s Atomic Orbitals (pay attention to size, shape, and number of each orbital type) Principal quantum numbers ( ) Principal energy levels Energy sublevels Table 5-2 Hydrogen s First Four Principal Energy Levels Principal quantum number (n) Sublevels (types of orbitals) present Number of orbitals related to sublevel Total number of orbitals related to principal energy level (n 2 )

13 Read pg Electron Configurations Ground State Electron Configurations Electron configuration The aufbau principle The Pauli exclusion principle Hund s Rule Orbital Diagrams and Electron Configuration Notations Orbital Diagram Electron configuration Note: 2p 3 is the same as Noble-gas notation 13

14 Exceptions to predicted configurations Ex Writing electron configurations Germanium (Ge), a semiconducting element, is commonly used in the manufacture of computer ships. What is the ground-state electron configuration for an atom of germanium? Valence electrons S [Ne]3s 2 3p 4 Cs [Xe] 6s 1 Fr [Rn] 7s 1 Electron-dot structures Table 5-5 Element Atomic Number Lithium 3 Beryllium 4 Boron 5 Carbon 6 Nitrogen 7 Oxygen 8 Fluorine 9 Neon 10 Electron configuration Electron-Dot Structure 14

15 This page for In Class Use Electromagnetic Radiation λ wavelength υ -frequency Electromagnetic Spectrum Quantized Energy Albert Einstein concluded that electromagnetic radiation behaves like waves & particles Max Plank concluded energy is quantized photon - Quantized energy means energy can only be gained or lost in specific increments. ELECTRONS ALSO BEHAVE LIKE BOTH WAVES AND PARTICLES. Atomic Emission Spectra Flame Tests Results: Compound Lithium Chloride Sodium Chloride Potassium Chloride Calcium Chloride Strontium Chloride Copper (II) chloride Flame Color Bohr s Model of the Atom 1) He suggested that the electron moves around the nucleus in only certain allowed. 2) He proposed that the atom has only certain allowable energy states. Absorption ground state lowest energy state (electrons in lowest possible energy orbitals) excited state electrons have jumped up to higher energy orbitals (gained energy) ionized state electrons have gained enough energy to be removed from the atom. Emission 15

16 3) Bohr s model was able to predict the frequencies of the lines in hydrogen s atomic emission spectrum. 4) Bohr s Model only worked in determining the atomic spectrum of hydrogen. n = 5 n = 4 n = 3 n = 2 n = 1 Quantum Mechanical Model of the Atom Heisenberg Uncertainty Principle Schrodinger Wave Equation Electrons are treated as, the equation can tell you the coordinates most probable to find an electron of a given energy. atomic orbital 16

17 Energy Level # electrons # orbitals needs Types & # s of orbitals Hydrogen only All polyelectronic atoms auf bau principle an electron will occupy the that can receive it. Pauli exclusion principle (no two electrons can have the same set of four quantum numbers) i.e. If two electrons are in the same orbital, they must have. Hund s Rule the lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals. i.e. as electrons fill into degenerate orbitals one must go into each orbital will spins before you. 17