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Chemistry Notes Week 4

by: Alyssa Anderson

Chemistry Notes Week 4 CHEM 1030

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Alyssa Anderson

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These notes cover what we covered on 2/2/16 and 2/4/16. They include a detailed periodic table and diagrams to aid in understanding the electron configuration.
Fundamentals Chemistry I
Dr. Streit
Class Notes
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This 9 page Class Notes was uploaded by Alyssa Anderson on Thursday February 4, 2016. The Class Notes belongs to CHEM 1030 at a university taught by Dr. Streit in Spring 2016. Since its upload, it has received 25 views.


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Date Created: 02/04/16
1 Chemistry Notes- Week 4 Quantum Theory 1. Atoms and photons at the microscopic level do not measure equally to the microscopic level 2. The laws for macroscopic are not applicable to the microscopic level 3. All energy is transferred through the measure of waves (E = h x v) 4. E = h x v means energy is calculated by multiplying Planks Constant (6.63 x 10^-34 J x s) by the frequency The Schrodinger Equation 1. Erwin Schrodinger realized the wave and particle characteristics were different in electrons 2. Particle behavior is determined by mass (m) while wave behavior is determined by the wave function (Ψ) in the equation H (m) x Ψ = E x Ψ 3. Quantum Mechanics- defines the region where electron is most likely to be at a given time 4. The probability of finding an electron in a certain area of space is proportional to Ψ^2 and is called electron density 5. Energy states and wave functions are characterized by a set of quantum numbers 6. Quantum numbers and wave functions describe atomic orbitals Quantum Numbers 1. They are required to describe the distribution of electron density in an atom 2. In order to describe an atomic orbital, you must know the three quantum numbers 2 3. Principal Quantum Number (n) A. Designates SIZE of the orbital B. The larger the value of n the larger the orbitals C. The allowed values of n are integral numbers (1, 2, 3, etc.) D. The collection of orbitals with the sam value of n are frequently called shells 4. Angular Momentum Quantum Numbers (l ) A. Describes the SHAPE of the orbital B. Values of l are integers that depend on the value of the principle quantum number n C. Allowed values of l range from 0 to n-1 D. The collection is called a subshell 5. Magnetic Quantum Number (ml ) A. Determines the ORIENTATION of orbitals in space B. Values of ml are integers that depend on the value of the angular momentum number l C. -l , 0, + l D. Quantum Numbers designate shells, subshells, and orbitals E. REFER TO TABLE 3.2 6. Speed (ms) A. It is not derived by the equation B. Found through experiments that included a beam of atoms that were split by a magnetic field C. It was concluded that electrons behave like tiny magnets D. Specifies the electrons spin E. ms = +/- 1/2 3 SUMMARY 1. Principle(n)- SIZE 2. Angular (l )- SLOPE/SHAPE 3. Magnetic (ml )- ORIENTATION Example: 2p^2 1. n = 2 2. l = p = 1 3. m l = +1, 0, -1 4. ms = +1/2, -1/2 Atomic Orbitals 1. All s orbitals are spherical in shape but alter in size ( 1s < 2s < 3s) 2. All p orbitals are dumbbell shaped and have 3 orientations 3. D orbitals vary and have 5 orientations 5. F orbitals vary and have 7 orientations 6. Energy of orbitals- in a hydrogen atom, depends only on n Aufbau Principle 1. States that electrons are added to the lowest energy orbitals first before moving to higher energy orbitals 2. Example: Li has 3 electrons so its configuration is 1s^2 / 2s^1 3. Example: Be has 4 electrons so its configuration is 1s^2 / 2s^2 4. Example: B has 5 electrons so its configuration is 1s^2 / 2s^2 / 2p^1 5. Example: C has 6 electrons so its configuration is 1s^2 / 2s^2 / 2p^2 6. Example: F has 9 electrons so its configuration is 1s^2 / 2s^2 / 2p^5 7. NOTE: 2p orbitals are degenerate 4 Pauli Exclusion Principle 1. No two electrons in an atom can have the same four quantum numbers 2. The principle number, angular momentum number, magnetic number, and speed cannot ALL be the same 3. Only 2 electrons can occupy an atomic orbital Hund’s Rule 1. The most stable arrangement of electrons is the one in which the number of electrons with the same spin is maximized 2. In other words, put 1 electron in each box before pairing NOTE: All the chemical and physical properties of matter are given by how the electrons are arranged in each orbital. Paramagnetism is when there are one or more unpaired electrons in an atom (such as the case of O and F). Diamagnetism is when all the electrons in an atom are paired, such as Neon. Electron Configuration 1. Describes where the electrons are distributed in the various atomic orbitals 2. In the ground state of hydrogen, the electron is found in the 1s orbital (1s^1 means the principal number n = 1 and the angular momentum is s = 0). If hydrogens electrons were found in a higher energy we would say the atom is in an excited state (2s^1) 3. In multi-electron atoms, the orientations of the orbitals are SPLIT (i.e. goes from 3s to 3p then from 4s to 3d) 5 Rules of Electron Configuration 1. Electrons will reside in the available orbitals of the lowest possible energy 2. Each orbital can accommodate a maximum of two electrons 3. Electrons will not pair in degenerate orbitals if an empty orbital is available 4. Orbitals will fill in the order indicated in the figure to the right. Worked example 3.10 Problem: What’s the electron configuration and orbital diagram of Ca (Z= 20)? Solution: 1s^2 / 2s^2 / 2p^6 / 3s^2 / 3p^6 / 4s^2 Noble Gas Core 1. The electron configurations of all elements except H and He can be represented by using a noble gas core 2. K (Z =19) has the configuration 1s^2 / 2s^2 / 2p^6 / 3s^2 / 3p^6 / 4s^1 but since argon (Ar) is 1s^2 / 2s^2 / 2p^6 / 3s^2 / 3p^6 you can adjust and only write [Ar] 4s^1 Electron configuration and the periodic table 1. Valence electrons are the outer electron involved in chemical reactions and can be identified by the period number 2. 4f = the lanthanide (rare earth) series 3. 5f = the actinide series 6 4. Notable exceptions to electron filling in the transition metals: A. Chromium (Z = 24) is [Ar] s3^1 / 3d^5 B. Copper (z = 29) is [Ar] 4s^1 / 3d^10 C. The reason for these anomalies is the slightly greater stability of d subshells that are either half filled (d^5) or completely filled (d^10) Discoveries in the Periodic Table A. In 1864 John Newlands noted that when the elements were arranged in order of atomic number, every eighth element had similar properties. They could be grouped according to their properties and he called it the law of octaves. B. In 1869 Dmitri Mendeleev and Lothar Meyer independently proposed they idea of periodicity. 1. Mendeleev grouped the 66 known elements according to their properties and atomic mass 2. Mendeleev predicted properties for elements not yet discovered such as gallium (Ga) 3. However, Mendeleev could not explain inconsistencies such as argon coming before potassium in the periodic table despite having a higher atomic mass 4. In 1913 Henry Mosley discovered the correlation between the number of protons (atomic number) and frequency of x-rays generated. 5. By ordering the periodic table by atomic number instead of atomic mass, scientist were able to make sense of discrepancies. 6. Entries today include atomic number and symbol and are arranged according to electron configuration 7 For more help on the periodic table check out! 8 Works Cited sa=i&rct=j&q=&esrc=s&source=images&cd=&cad=rja&uact=8&ved=0ahUKEwi C0-n04N7KAhULbD4KHYJqC8QQjB0IBg& %2Facademy%2Flesson%2Fground-state-electron-configuration-definition- example- quiz.html&psig=AFQjCNF8UKlnbc139XTFXnGPJGogs6GOtw&ust=1454697555 658620 Electronic_Structure_of_Atoms_and_Molecules/ Electronic_Configurations 9 Exceptions to the Rule: Chromium: Z = 24 —> [Ar] 4s^1 / 3d^5 Copper: Z = 29 —> [Ar] 4s^1/ 3d^10 Niobium: Z = 41 —>[Kr]4d^4 / 5s^1 Molybdenum: Z = 42 —> [Kr] 4d^5 / 5s^1   Ruthenium: Z = 44 —> [Kr] 4d^7 / 5s^1   Rhodium:  Z = 45 —> [Kr] 4d^8/ 5s^1   Palladium: Z = 46 —> [Kr] 4d^10   Silver: Z = 47 —> [Kr] 5s^1 / 4d^10 Lanthanum: Z = 57 —> [Xe] 6s^2 / 5d1 Actinium: Z = 89 —> [Rn] 7s^2 / 6d^1 Cerium: Z = 58 —> [Xe] 6s^2 / 4f^1 / 5d^1 Thorium: Z = 90 —> [Rn] 7s^2 / 6d^2 Gadolinium: Z = 64 —> [Xe] 6s^2 / 4f^7 / 5d^1 Protactium: Z = 91 —> [Rn] 7s^2 / 5f^2 / 6d^1 Platinum: Z = 78 —> [Xe] 6s^1 / 4f^14 / 5d^9 Uranium: Z = 92 —> [Rn] 7s^2 / 5f^3 / 6d^1 Gold: Z = 79 —> [Xe] 6s^1 / 4f^14 / 5d^10 Neptunium: Z = 93 —> [Rn] 7s^2 / 5f^4 / 6d^1   Curium: Z = 96 —> [Rn] 7s^2 / 5f^7 / 6d^1   Lawrencium: Z = 103 —> [Rn] 7s^2 / 5f^14 / 7p1


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