CHEMISTRY CHAPTER 5 OUTLINE NOTES 5.1 – Light and Quantized Energy • The Nuclear Atom and Unanswered Questions o Bohr and the Planetary Model • Wave Nature of Light & Spectroscopy o Absorption of Energy and the “Excited” State o Electromagnetic Radiation o Wavelength and Frequency o Speed of Light Equation: c = λν o Electromagnetic Spectrum: Fig 5-5 pg 120 o Visible Light : ROYGBIV* o Practice Problems: pg 121 #’s 1-4 • Particle Nature of Light o Quantum Concept o Energy of Light Equation o Photoelectric Effect o Practice Problems #’s 5 and 6 on pg 124 • Atomic Emissions Spectra • Atomic Absorption Spectra* 5.2 – Quantum Theory and the Atom • Bohr Model of the Atom • The Quantum Mechanical Model of the Atom • The Heisenberg Uncertainty Principle • Hydrogen’s Atomic Orbitals 5.3 – Electron Configurations • Ground State Electron Configurations • Orbital Diagrams and Electron Configuration Notation • Valence Electrons and Lewis Dot Structures • Practice Problems

CHEMISTRY CHAPTER 5 NOTES 5.1 – Light and Quantized Energy • The Nuclear Atom and Unanswered Questions o Although Rutherford’s scientific model of an atom was a breakthrough, it lacked detail about how electrons occupy the space surrounding the nucleus of an atom. o Questions Still Unanswered: • How are an atom’s electrons arranged in space around the nucleus? • Why aren’t negatively charged electrons pulled in towards the nucleus? • Why are there differences in chemical behavior among the various elements? o One scientist named Neils Bohr thought of electrons being in “orbit” around the nucleus in much the same manner as Earth is in orbit around the sun. This is sometimes called the planetary atomic model or Bohr’s Model. o So a hydrogen atom should be similar to a solar system consisting of a sun and one planet. o Scientists began to unravel the puzzle of chemical behavior in the early 1900’s and observed that certain elements emitted visible light when heated in a flame. o Analysis of the emitted light revealed that an elements chemical behavior is related to the arrangement of the electrons in its atoms. • Wave Nature of Light & Spectroscopy o When a substance is exposed to a certain intensity of light or some other form of energy, the atoms absorb some of that energy and are said to be in an “excited” state. o Spectroscopy is the study of the interaction of matter and radiant energy. It is a method of studying substances that are exposed to some sort of exciting energy. o Electromagnetic Radiation: A form of energy that exhibits wavelike behavior as it travels through space. • Types of electromagnetic radiation include waves ranging from very low to very high energy (a product of frequency). Visible light is a component of this spectrum in addition to radio waves, ultraviolet light, infrared light, x-rays, gamma rays. o Wavelength and Frequency • Wavelength (λ) a.k.a. lambda = the distance between two similar points on two successive waves. Measured in meters. • Amplitude = the height the crest or depth of the trough. Refers to the intensity or brightness of a light. • Frequency (ν) a.k.a. nu= the number of waves that pass a point per second. This is what changes the energy of a wave. Hertz (Hz) is a measure of frequency or cycles. We commonly use MHz or megahertz (1 MHz = 1 x 106 Hz).

o Speed of Light Equation: c = λν • The product of frequency (in Hz or s-1) and wavelength (in meters) is equal to the speed of light which is a constant. • Speed of Light: c = 3.00 x 108 m/s o Electromagnetic Spectrum • Copy Fig 5-5 on pg 120 into your notes.

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o Visible Light : ROYGBIV* • Red (lowest energy) • Orange • Yellow • Green • Blue • Indigo • Violet (highest energy) o Practice Problems: pg 121 #’s 1-4 Particle Nature of Light – The wave theory of light does not explain its everyday behavior and fails to adequately describe important aspects of lights interactions with matter. o Quantum Concept • Max Planck began searching for explanation for light emitted by heated objects and he found that matter can gain or lose energy only in small, specific amounts called quanta.

o Energy of Light Equation: E = hv • E is energy measured in joules (J) • h is Planck’s constant and is equal to 6.626 x 10-34 J.s • v is frequency in 1/s o Photoelectric Effect • The process by which electrons are ejected from a substance when light of a certain frequency shines on its surface. This type of technology is used in solar powered devices such as calculators

o Practice Problems #’s 5 and 6 on pg 124

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Atomic Emissions Spectra is a set of frequencies of the electromagnetic waves emitted by atoms of the element. These spectrums appear as colored lines and are unique for each individual element. They can be used to identify an element in much the same way a fingerprint identifies a person. o Lithium – red o Sodium – yellow o Potassium – violet o Calcium – red/orange o Strontium – bright red o See Figure 5-9 on pg 126 of your textbook.

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Atomic Absorption Spectra is produced when white light is passed through a sample and certain wavelengths of light are absorbed which then produces dark lines on a continuous spectrum.

5.2 – Quantum Theory and the Atom • Bohr Model of the Atom o According to his model, hydrogen’s atomic emission spectrum results from electrons dropping from higher-energy atomic orbits to lower-energy atomic orbits. • The Quantum Mechanical Model of the Atom o It is based on the assumption that electrons are waves. • The Heisenberg Uncertainty Principle o It is impossible to know both the speed (velocity) and position of an electron. You may only know one or the other at one time. • Hydrogen’s Atomic Orbitals o Electrons occupy three-dimensional regions of space called atomic orbitals. o Schrodinger’s Wave equation can predict the probable location of the electrons in an atom. o There are 4 orbital types based on energy sublevel: • s – sphere shaped, 1 on each energy sublevel • p – lobe shaped, 3 on each energy sublevel • d – complex shaped, 5 on each energy sublevel • f – complex shaped, 7 on each energy sublevel

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Quantum Numbers: o n = principle energy level (1-7) o l = sublevel (s, p, d and f) o m = orbital (2 e- max per orbital and shape related to sublevel) o s = spin (up or down)

5.3 – Electron Configurations • Ground State Electron Configurations o Aufbau Principle – electrons fill lowest energy levels first. o Pauli Exclusion Principle – only 2 electrons can occupy an orbital and must have opposite spins (means they move in opposite directions) o Hund’s Rule – electrons fill the orbitals of an energy level with the same spin first then the opposite spin. • Orbital Diagrams and Electron Configuration Notation Filling sequence:

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Practice Problems: o Pg 139 #’s 18-22

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Valence Electrons and Lewis Dot Structures o The outer most electrons (s and p orbitals only) of an atom are called the valence electrons. These are the electrons involved in chemical reactions. o It is often useful to draw these outer electrons around the symbol of an element. This notation is called a Lewis Electron Dot Diagram. Practice Problems: o Pg 141 # 23

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o Write out the electron configuration, orbital filling diagram and dot structure for elements with atomic numbers 1-20 plus Ni, Ce and Sn. You may skip any elements done together in class as examples in your notes.