Honors Unit 6 Notes - Atomic Structure

Transcription

1 Name: Honors Unit 6 Notes - Atomic Structure Objectives: 1. Students will have a general understanding of the wave nature of light and the interrelationship between frequency, wavelength, and speed of electromagnetic radiation. 2. Students will have a general understanding of the quantum hypothesis, Einstein s photoelectric effect, and Bohr s contributions to the model of the atom. Based on modern theories of the behavior of electrons, students will be able to explain the production of bright line spectra and will have a basic understanding of the particlewave nature of matter. 3. Students will be able to represent electrons in terms of electron configurations, orbital notations, and core notations for ground state atoms, excited atoms, and ions. 4. Students will develop an understanding of the relationship between electron configuration and the structure of the periodic table. Students will begin to develop and understanding of the relationship between electron configuration and various properties of individual atoms: magnetic properties, atomic radii, ionic radii, ionization energy, electron affinity, and metallic, nonmetallic and noble gas properties. Light, Photon Energies, and Atomic Spectra [pgs in textbook] o Electromagnetic waves o Visible light a form of electromagnetic radiation that is perceivable to human beings and is seen in the colors of the rainbow

2 2 Wave Vocabulary o Crest o Trough o (λ ) the distance from crest to crest or trough to trough in a wave Units: o (ν ) the number of wavelengths that pass a given point in a second Units: o - the distance from the origin to the crest or the trough of a wave o (represented by the variable ) the rate at which all forms of electromagnetic radiation travel through a vacuum Speed of light = Relationship between Wavelength & Frequency As Wavelength increases, frequency. As Wavelength decreases, frequency.

3 3 Wave Equation (Wave Theory of Light) One equation relates speed, frequency, and wavelength: o c = o λ = o ν = Example #1: The wavelength of the radiation which produces the yellow color of sodium vapor light is nm. What is the frequency of this radiation? The Electromagnetic Spectrum o Complete range of wavelengths and frequencies o Mostly invisible

4 o The visible spectrum continuous spectrum; components of white light split into its colors ROY G. BIV From (violet) to (red) Can be split by a prism 4 Particle Theory of Light o Light is generated as a stream of light particles called. o Equation: E = h = ν = Example #1: (a) If the frequency of a ray of light is 5.09 x Hz, calculate the energy, in Joules, of a photon emitted by an excited sodium atom. (b) Calculate the energy, in kilojoules, of a mole of excited sodium atoms. Example #2: What is the energy of a photon from the green portion of the rainbow if it has a wavelength of 4.90 x 10-7 m?

5 5 Bohr Model of the Atom [pgs in textbook] When an electron absorbs a photon of energy, the electron jumps from its ground state to its excited state. o Ground state o Excited state o Line Spectra (a.k.a Atomic Emission Spectra) Unique for every element Used to identify unknown elements To emit light, an electron will drop from its excited state to its ground state. As it falls, energy is emitted that we see as light. Is Light a Particle or Wave? o (1921) Albert Einstein wins Novel Prize in Physics for the photoelectric effect o Photoelectric effect occurs when strikes the surface of a metal and are ejected. o Conclusion: Light not only has but also has properties. These mass-less particles, called, are packets of energy. Light has a!

6 6 Quantum Mechanics [pgs in textbook] o Quantum mechanics o Erwin Schrödinger wave equation that describes hydrogen atom o - the exact location and speed of an electron cannot be determined simultaneously (if you try to observe it, you interfere with the particle) You can know either the location or the velocity, but not both! Electrons exist in and not on specific rings or orbits like in the Bohr model of the atom. Atomic Structure Quantum numbers o They range from the most general locator to the most specific. 1. Energy Level (n) o Always a positive integer o Indicates size of orbital, or how far is from Larger n value = o Similar to Bohr s energy levels or shells o n = for given element 2. Sublevel o Indicates of orbital o Letters

7 7 Energy level 1 has Energy level 2 has Energy level 3 has Energy levels 4-7 have 3. Orbital o The most specific piece of information is about the number and location of the electrons within the sublevel The sublevel has The sublevel has The sublevel has The sublevel has o Orbital Every orbital can hold! Shapes of atomic orbitals: s = p = d = f =

8 8 Capacities of Levels, Sublevels, and Orbitals Principal Energy level (n) 1 Sublevels Present (s, p, d, or f) Number of Orbitals Present s p d f Total Number of Orbitals Maximum Number of Electrons in Energy Level Rules for how Electrons fill into the Electron Cloud Electrons take position into the cloud according to a set of rules Aufbau Principle Pauli Exclusion Principle Hund s Rule An Introduction to Electron Configuration [pgs in textbook] The system of numbers and letters that designates the most probable locations of e - 3 major methods: o o o

9 9 Full Electron Configuration Example Notation: 1s 2 2s 1 Pronounced: A. What does the coefficient mean? B. What does the letter mean? C. What does the exponent mean? Steps for Writing Full Electron Configurations: 1. Determine the total number of electrons the atom has (for neutral atoms it is equal to the atomic number for the element). Example: F atomic # = # of p+ = # of e- = 2. Fill orbitals in order of increasing energy. 3. Make sure the total number of electrons in the electron configuration equals the atomic number. Aufbau Chart (Filling Order of Energy Levels) When writing electron configurations: d sublevels are n 1 from the row they appear in on the periodic table f sublevels are n 2 from the row they appear in on the periodic table

10 Practice! Write full electron configurations for the following elements: 10 Nitrogen: Helium: Phosphorous: Rhodium: Bromine: Cerium: Noble Gas/Abbreviated Configuration A. Where are the noble gases on the periodic table? B. Why are the noble gases special? C. How can we use noble gases to shorten regular electron configurations? 1. Look at the periodic table and find the noble gas in the row above where the element is. 2. Start the configuration with the symbol for that noble gas in brackets, followed by the rest of the electron configuration. Example: Tin

11 Practice! Write noble gas configurations (abbreviated electron configurations) for the following elements 11 Sulfur: Rubidium: Bismuth: Zirconium: Orbital Diagrams (Orbital Diagram Configuration) Orbital diagrams use (sometimes circles) to represent and. are used to represent the. = Sublevels = Don t forget - orbitals have a capacity of!! Two electrons in the same orbital must have so draw the arrows pointing in. Example: Orbital diagram for oxygen

12 Drawing Orbital Diagrams 1. First, determine the electron configuration for the element. 2. Next draw boxes for each of the orbitals present in the electron configuration. Boxes should be drawn in order of increasing energy (see the Aufbau chart). 3. Arrows are drawn in the boxes starting from the lowest energy sublevel and working up. This is known as the Aufbau principle. Add electrons one at a time to each orbital in a sublevel before pairing them up (Hund s rule) The first arrow in an orbital should point up; the second arrow should point down (Pauli exclusion principle) 4. Double check your work to make sure the number of arrows in your diagram is equal to the total number of electrons in the atom. # of electrons = atomic number for an atom 12 Practice! Draw orbital diagrams for the following elements: Nitrogen: Nickel: Electron Configurations: Periodic Relationships [pgs in textbook] o Valence electrons o They determine the chemical properties of an element.

13 o Write the noble gas configuration for an element the valence electrons are the ones 13 Electron Configurations of Ions Cations: (Z = 11) Na Draw Na + Anions: Draw Cl - Transition metals: lose electrons from the highest principal energy level (n) first, then they lose their d electrons Electron configuration for Zr atom = Electron configuration for Zr +2 =

14 Periodic Properties &Trends [pgs in textbook] 1. Electronegativity (See page 8 in your reference book for values on the Pauling scale.) ***TREND: Increases going and to the. Across a period Down a group 2. Atomic Radius ***TREND: Increases going and to the. Down a group Across a period ***Remember*** LLLL Lower, Left, Large, Loose

15 3. Ionic Radius 15 o Ionic Radius of Cations o Ionic Radius of Anions ***Cations are than the atoms from which they form. ***Anions are than the atoms from which they form. Trends in ion sizes are the same as the trends in atom sizes. 4. Ionization Energy o 1 st I.E. = o 2 nd I.E. = o 3 rd I.E. = Ex. B --> B + + e- Ex. B + --> B +2 + e- Ex. B +2 --> B +3 + e- I.E. = 801 kj/mol I.E.2 = 2427 kj/mol I.E.3 = 3660 kj/mol ***TREND: Increases going and to the. Down a group Across a period ***Remember*** LLLL Lower, Left, Large, Loose

16 5. Metallic Character 16 ***TREND: Increases going and to the. Think about where the metals and nonmetals are located on the periodic table to help you remember the trend for metallic character!! 6. Electron Affinity ***TREND: Increases going and to the. Summary of Periodic Trends: