General Chemistry Notes Week 3
General Chemistry Notes Week 3 Chem 1315-003
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Popular in Chemistry
This 10 page Class Notes was uploaded by Ryan Henry on Sunday February 7, 2016. The Class Notes belongs to Chem 1315-003 at University of Oklahoma taught by Dr. Awasabisah in Spring 2016. Since its upload, it has received 11 views. For similar materials see General chemistry in Chemistry at University of Oklahoma.
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Date Created: 02/07/16
General Chemistry Notes Week 3 Review pictures: (Note that radio waves aharmless, while gamma rays are the mos owerful AND dangerous form of radiation EX: the radiation emitted by an atomic bomb, chernobyl, and fallout 4) (This is an example of diffraction in a light source, makes is easier to understand) The wave behavior of matter ● Proof that the electron has wave nature came a few years later with the demonstration that a beam of electrons would produce an interference pattern the same as waves do0 ● . ● De Broglie’s equation allows the calculation of the wavelength for a particle, given the mass and velocity. ● de Broglie’s equation also allows the calculation of the velocity of a particle, given the mass and wavelength. ● Uncertainty Principle: ○ Heisenberg showed that the more precisely the velocity of a particle is known, the less precisely its position is known and vice versa: ○ ○ Complementarity of position (particle nature) and velocity (wave nature). It is not possible to observethe interference pattern of 1 electron and simultaneously determine which slit it went through. ○ An electron is observed as either a particle or a wave. Not both! ○ Shows the probability of an electron behaving a certain way! Quantum mechanics and atomic orbitals ● Erwin Schrödinger developed a mathematical model, known as quantum mechanics, into which both the wave and particle nature of matter could be incorporated: ● Shows the probability of an electron behaving a certain way! Quantum Numbers ● Schrödinger’s wave equation yields a set of wave functions, or orbitals, and their corresponding energies. ○ Each orbital describes a spatial distribution of electron density. ● An orbital is not the same as an orbit –the motion of the electron in an atom cannot be tracked precisely. ● The Bohr modelsuggested a single quantum number, n, to describe an orbit. ● The quantum mechanical modeluses three quantum numbers to describe the atomic orbital (n, l, and ml ). ● Use four quantum numbers to describe the electron (n, l, and ml , ms) The 4 Quantum numbers: 1. Principal Quantum Number, n Designates the size of an orbital a. The principal quantum number, n, describes the energy level on which the orbital resides. b. The values of n are integers ≥ 1. (n= 1, 2, 3, etc.) c. n designates the size of the orbital the larger the value of n, the farther the orbital is from the nucleus. 2. Azimuthal quantum number (or azimuthal quantum number) l Describes the shape of orbital a. This quantum number defines the shape of the orbital. b. Allowed values of are integers ranging from 0 to n − 1. c. Letter designations are used to communicate the different values of land, therefore, the shapes and types of orbitals. (spdf = 0123) d. EX: if n=1 then it is s, if n=2 then it is (s or p), and so on 3. Magnetic quantum number, ml Describes the orientation of orbital in space a. This quantum number describes the threedimensional orientation of the orbital. b. Values are integers ranging from lto l: −l≤ ml≤ l. c. On any given energy level, there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals. 4. Electron Spin Quantum Number, ms Describes the Spin of an electron that occupies an orbital a. In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy. b. The “spin” of an electron describes its magnetic field, which affects its energy. c. This discovery led to a fourth quantum number, the spin quantum number, ms. i. The spin quantum number has only 2 allowed values: +1/2 and −1/2. ● An orbital can hold a maximum of two electrons. ● Pauli Exclusion Principle: No two electrons in the same atom can have the same four quantum numbers. ● The total number of orbitals in a shell is n2 Atomic orbitals ● s Orbitals: ○ Value of l= 0. ○ Spherical in shape ○ Radius of sphere increases with increasing value of n. ○ All s orbitals have the same shape but are different in size. ○ s orbitals possess n−1 nodes, or regions where there is 0 probability of finding an electron. ○ ● p Orbitals ○ Value of l= 1. ○ Have two lobes with a node between them. ○ Subscripts denote the axis along with which the three p orbitals lie (x, y, and z) ○ ● d Orbitals ○ Value of lis 2. ○ Four of the five orbitals have 4 lobes (“clover leaf”); the other resembles a p orbital with a doughnut around the center. ○ ● f Orbitals ○ The seven possible 4f orbitals (8 lobes and 3 nodal planes). ○ The Periodic Table ● There are three main categories on the periodic table: ○ Metalselements that tend to lose electrons during chemical change, forming positive ions. ○ Nonmetalssubstances whose atoms tend to gain electrons during chemical change, forming negative ions. ○ Metalloids(semimetals) have properties intermediate between metals and nonmetals. ■ Metalloids form a narrow diagonal band between metals and nonmetals (Al is an exception –classified as a metal). ■ The properties of metalloids fall between those of metals and nonmetals. ● Group or family: vertical (up and down) column of elements that have similar chemical properties. ● Period: horizontal (side to side) row in the periodic table. ● The Periodic Table: Alkaline Earth Metals ○ They are fairly reactive, but not quite as reactive as the alkali metals. ● The Periodic Table: Halogens ○ They are always found in nature as a salt. Electron Configurations ● Electron configurations describe the distribution of all electrons in an atom. ● Orbitalsfill in order of increasing energy, with no more than two electrons per orbital. ○ The ground state of an atom is the most stable electron configuration of an atom. ● Orbital diagrams represent electrons with halfarrows and orbitals with boxes or lines. ○ The direction of the arrow represents the spin of the electron. Aufbau Principle : Electrons fill the lowest energy orbitals first. Pauli Exclusion Principle : No two electrons in an atom may have the same set of four quantum numbers.Therefore, no orbital may have more than two electrons, and they must have opposite spins. Hund’s Rule: For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized. This means that electrons will not pair up if an empty orbital at the same energy level (same n, same l) is available. Condensed Electron Configurations: ● The outer shell electrons are called the valence electrons ● Inner shell electrons are called core electrons ● Electron configurations may represent all electrons, or may emphasize the valence electrons by using condensed electron configurations (For elements with a high amount of electrons, the condensed electron configuration simplifies the amount of electrons written) EX: 1s22s22p63s23p64s23d104p65s24d105p2 => [Kr] 5s24d105p2 Anomalies: ● Irregularities occur when there are enough electrons to halffill sand d orbitalson a given row. ● Chromium: [Ar] 4s1 3d5 rather than the expected [Ar] 4s2 3d4 ● Because of sublevel splitting, the 4s sublevel is lower in energy than the 3d; therefore, the 4s fills before the 3d. ○ The difference in energy is not large ● Some of the transition metals have irregular electron configurations in which the ns only partially fills before the (n−1)d or doesn’t fill at all. ○ Therefore, their electron configuration must be found experimentally. Coulomb’s law describes the attractions and repulsions between charged particles. ● For like charges, the potential energy (PE) is positive ● For opposite charges, the potential energy is negative ● Shielding: The amount of repulsion an electron experiences from other electrons.
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