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CHEM 1030 Cagg Chapter 3.5-3.10 Notes

by: Amy Notetaker

CHEM 1030 Cagg Chapter 3.5-3.10 Notes Chem 1030

Marketplace > Auburn University > Chemistry > Chem 1030 > CHEM 1030 Cagg Chapter 3 5 3 10 Notes
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These are lecture/book notes for the material in chapter 3.5-3.10.
Fundamental Chemistry I
Brett A Cagg
Class Notes
Cagg, Chemistry
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This 4 page Class Notes was uploaded by Amy Notetaker on Sunday February 21, 2016. The Class Notes belongs to Chem 1030 at Auburn University taught by Brett A Cagg in Spring 2016. Since its upload, it has received 42 views. For similar materials see Fundamental Chemistry I in Chemistry at Auburn University.


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Date Created: 02/21/16
Lecture/Book Notes: Chapter 3.5-3.10 CHEM 1030 Cagg Highlighted: Vocab ----- Highlighted: Formula Section 3.5 v The de Broglie Hypothesis • De Broglie claimed that an electron in an atom behaves like a standing wave, and if it does so in a hydrogen atom, the wavelength should fit the circumference of the orbit exactly. • To show the relationship between circumference if an allowed orbit and the wavelength of the electron, use this formula: 2πr = nλ - “r” is the radius of the orbit - “λ" is the wavelength of the electron wave - “n” is the positive integer • He later concluded that waves behave like particles and particles have wave like ▯ properties, wave properties and particles are related by this formula: λ = ▯▯ v Diffraction of Electrons • Experiments conducted by 3 physicists who demonstrated and verified that electrons do have wave like properties. Section 3.6 • Although there were theories to explain where electrons were located, scientists were still confused. • They also did not fully understand the behavior of electrons. v The Uncertainty Principle • The Heisenberg Uncertainty Principle: states that it is impossible to know both the momentum (mass × velocity) and the position of x with certainty. This is the ▯ ▯ formula which represents this principle: ∆x × ∆p ≥ ▯▯ or ∆x × m∆u ≥ ▯▯ • This principle is applied to a hydrogen atom, and it is concluded that the electron cannot orbit the nucleus in a well-defined path. v The Schrodinger Equation • An Austrian physicist formulated an equation, which describes the behavior and energies of submicroscopic particles. • The Schrodinger Equation: incorporates the particle behavior in terms of “m” (mass) and wave behavior in terms of ψ (psi), which depends on the location in space of the system. • This equation developed a new field called quantum mechanics/wave mechanics v The Quantum Mechanical Description of the Hydrogen Atom • Quantum mechanics does not specify the exact location of an electron in an atom, but it does tell us the approximate region where the electron will most likely be at a certain time. • Electron density: a concept, which gives the probability that, an electron will be found in a certain part of an atom. • Atomic orbital: the wave function of an electron in an atom. Section 3.7 • There are 3 quantum numbers, which describe the distribution of electron density in an atom: principal quantum number, angular momentum quantum number, and the magnetic quantum number. v Principal Quantum Number (n) • Principal quantum number: (n) tells you the size of the orbital. • The larger the “n” is, the greater the distance of the electron in the orbital to the nucleus is, which means the orbital is large as well. • The “n” can have integral values of 1, 2, 3, … • The value of “n” determines the energy of an orbital, but that is not true if the atom has more than 1 electron. v Angular Momentum Quantum Number (ℓ) • Angular Momentum Quantum Number: (ℓ) describes the shape of the atomic orbital. • The values of “ℓ” depend on the values of “n”, so for a value of “n” the possible values for “ℓ” range from 0 to n-1 - If n=1, then there is just 1 possible value for "ℓ" which is 0, because 1-1=0 - If n=2, then there are 2 possible values for "ℓ" which are 0 and 1, because 2-1=1, and the number before 1 is 0 - If n=3, then there are 3 possible values for "ℓ" which are 0, 1, and 2, because 3-1=2, and the numbers before 2 are 0 and 1 • The value of "ℓ" is designated by 4 letters: s, p, d, and f ℓ 0 1 2 3 Orbital s p d f designation • Shell: a collection of orbitals with the same value of “n” • Subshell: one or more orbitals with the same “n” and "ℓ" v Magnetic Quantum Number (m ) ℓ • Magnetic quantum number: (m ) desℓribes the orientation of the orbital • In a subshell, the “m” value depends on the "ℓ" value. For each "ℓ"value, there are 2ℓ + 1 values for "m ℓ - If ℓ=0, then there is just 1 possible value for ℓm " which is 0 - If ℓ=1, then there are 3 possible values for “mℓ" which are -1, 0, and +1 - If ℓ=2, then there are 5 possible values for “m " which are -2, -1, 0, +1, and ℓ +2 v Electron Spin Quantum Number (m ) ▯ • Electron spin quantum number: (m ) d▯scribes the electron’s spin. • There are 2 possible directions of spin +1/2 and -1/2 Section 3.8 • Orbitals don’t have a well-defined shape since the wave in the orbital extends from the nucleus to infinity. v S Orbitals • A s orbital is a spherical shape that is symmetric around the atom’s nucleus. • Happens when the “n” is 0. • As the energy levels increase, the electrons grow further from the nucleus. • 1s orbital: contains very little electrons. • 2s orbital: similar to the 1s orbital - Node: a space where there are no electrons and no possibility of finding any. • 3s orbital: larger than 1s and 2s orbitals, and contains 3 nodes. v P Orbitals • A p orbital points in a particular direction and is in the shape of 2 lobes. • Happens when the “n” is 2 or greater. • Also increases in size (2p, 3p, 4p, etc.). v D Orbitals • D orbitals have both complicated shapes and names. • Happens when “n” is 3 or greater. v F Orbitals • F orbitals have difficult shapes to represent. • These are important in accounting for the behaviors of elements with atomic numbers greater than 57. v Energies of Orbitals • Regardless of the subshell, orbitals in the same shell have the same energy. - This means that in the second shell, the orbitals, one 2s and three 2p have the same energy. - In the third shell, the orbitals, one 3s, three 3p, and five 3d have the same energy. - In the fourth shell, the orbitals, one 4s three 4p, five 4d and seven 4f have the same energy. Section 3.9 • Electron configuration: how the electrons are distributed in the atomic orbitals. v The Pauli Exclusion Principle • The Pauli exclusion principle: states that no two electrons in an atom can have the same 4 quantum numbers. v The Aufbau Principle • The Aufbau principle: makes it possible to build the periodic table to determine their electron configuration by steps. v Hund’s Rule • Hund’s rule: states that when the number of electrons in the same spin is maximized, that is when the arrangement of electrons in orbitals of equal energies is the most stable. • Diamagnetic: atoms with all paired electrons. • Paramagnetic: atoms with one or more unpaired electrons. v General Rules for Writing Electron Configurations • Electrons will be situated in the available orbitals with the lowest energy. • Each orbital can hold up to 2 electrons max. • If an empty orbital is available, electrons will not pair in the degenerate ones. • Orbitals will fill up in order. Section 3.10 • Noble gas core: another way an electron configuration can be shown of all elements except for hydrogen and helium.


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