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Molecularity CH1010 Week 3 Notes

by: Fabio Notetaker

Molecularity CH1010 Week 3 Notes CH1010

Marketplace > Worcester Polytechnic Institute > Chemistry > CH1010 > Molecularity CH1010 Week 3 Notes
Fabio Notetaker

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Covers the material from Tuesday, Thursday, and Friday's lectures 9/6/16~9/9/16.
Chem 1010 Molecularity
Prof. Burdette
Class Notes
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This 7 page Class Notes was uploaded by Fabio Notetaker on Friday September 9, 2016. The Class Notes belongs to CH1010 at Worcester Polytechnic Institute taught by Prof. Burdette in Fall 2016. Since its upload, it has received 25 views. For similar materials see Chem 1010 Molecularity in Chemistry at Worcester Polytechnic Institute.


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Date Created: 09/09/16
Molecularity CH1010 Week 3 Notes    Waves of Light:    Electromagnetic radiation​ is a form of energy characterized by waves that have an electric and  magnetic component, commonly known as l ​ ight.​     The ​electromagnetic spectrum​ depicts different types of EM radiation ordered by how much  energy each type contains.    A wave’s ​wavelength​ (represented by λ and measured in meters) is the distance between two  consecutive peaks of that wave.  A wave’s ​frequency ​(represented by ν and measured in s​  or  ­1​ Hertz) is the number of wavelengths that pass through a stationary point in one second.        Examples of high­energy (short wavelength, high frequency) EM waves include:  ● Gamma rays  ● X­rays  ● Ultraviolet rays    Examples of low­energy (long wavelength, low frequency) EM waves include:  ● Visible light  ● Infrared  ● Microwaves  ● Radio waves      The ​speed of light​ is represented by c and defined as 2.998 * 10​  m/s.  The speed of light can  obtained by multiplying an EM wave’s wavelength by its frequency:        Atomic Spectra:    Atomic emission spectra ​are the specific wavelengths of light produced by burning certain  elements.  ​Atomic absorption spectra ​are the specific wavelengths of light absorbed by certain  gases.  Each element has its own unique spectrum.    Examples of emission spectra:        Example of an absorption spectrum:      Particles of Light: Quantum Theory:    The energy released by EM waves is ​quantized​, meaning the energy is released in the form of  particles known as p ​ hotons.​     The amount of energy released by an EM wave can be obtained by multiplying its frequency by  Planck’s constant​, which is represented by h and defined as 6.626 * 10​  joule­seconds (J*s): 4​       Light can discharge an electron from the surface of a metal as long as the energy released by the  light is equal to or greater than the minimum amount of energy required to discharge the electron  (represented by Φ).  If the light’s energy is greater than necessary, the discharged electron will  have additional kinetic energy.  Φ can be obtained by subtracting the kinetic energy of the  electron from the energy released by the light:        Hydrogen Spectrum and the Bohr Model:    The ​Bohr model​ of the atom says that outside the nucleus, electrons can exist in different  energy levels​ in which the level closest to the nucleus is considered the first level:        The energy released from an electron moving from one energy level to another (represented by  ΔE) can be obtained by finding the difference of the inverse squares of both energy levels  (represented by n​  and n​ ) and multiplying by the constant ­2.178 * 10​  J:  ­18​ initial​ final​       This energy is as light.  Thus, for hydrogen in particular, the ​Rydberg equation​ can be used to  calculate the wavelengths of light emitted from hydrogen electrons moving from one energy  level to another:        R​H​is a constant defined as 1.097 * 10​  m​ .  7​ ­1​   Electrons as Waves:    Objects (light and electrons in particular), can behave as both particles and waves.  The  wavelength of a wave produced by an object can be obtained multiplying the object’s mass  (represented by m) and speed (represented by μ) and dividing Planck’s constant by that product:        Quantum Numbers:    Quantum numbers​ describe where certain electrons can be found outside an atom’s nucleus.    The ​principle quantum number​ (represented by n) can be any positive integer value (not  including 0) and represents an electron’s energy level.    The ​angular momentum quantum number​ (represented by ℓ) can range from 0 to n­1 and  represents the shape of the electron’s ​orbital​ (represented as s, p, d, or f).    The ​magnetic quantum number ​(represented by m​ ) can range from ­ℓ to +ℓ and represents the  ℓ​ electron’s orientation.    The ​spin quantum number ​(represented by m​ ) can be either posis​ve or negative ½ and  represents the direction of an electron’s spin.    Any set of one each of the first three quantum numbers define a unique orbital.  Adding a spin  quantum number to that set defines a unique electron.      Sizes and Shapes of Atomic Orbitals:    An atomic orbital describes the area in which certain electrons are likely to exist.  The four types  of orbitals (s, p, d, and f) have different shapes:          The Periodic Table and Filling Orbitals:    An atom’s ​electron configuration​ what types of orbitals that atom has and which energy levels  those orbitals are in.      For example, the electron configuration for lithium is 1s​ 2s​ .  The coefficient represents an  orbital’s energy level, the letter represents the type of orbital, and the superscript represents the  number of electrons in the orbital.  Adding another electron gives 1s​ 2s​  which is the  2​ 2​ 2​ 2​ 1​ configuration for beryllium.  Adding another electron gives 1s​ 2s​ 2p​  which is the configuration  for boron, and so on.    All the electrons in the outermost energy level are referred to as ​valence electrons​ and all the  other electrons are referred to as ​core electrons.​     Effective nuclear charge​ or ​Zeff​ is the net positive charge felt by an atom’s valence electrons.  This is calculated by subtracting the number of core electrons an atom has from the atom’s  atomic number.    For example, lithium has an atomic number of 3 along with 2 core electrons.  Subtracting the two  gives the effective nuclear charge, in this case +1.  In the cases of hydrogen or helium, the  effective nuclear charge is equal to the atomic number.    Orbitals are filled in order from lowest energy to highest energy according to the ​aufbau  principle​, and each orbital can contain up to two electrons:      This means that orbitals gets filled in pairs and according to the ​Pauli exclusion principle​, the  two electrons in a specific orbital will always have opposite spins.  The two spin directions are  represented by ⇃ and ↾ and two opposite are represented by ⇃↾.    In the case of p, d, or f orbitals, electrons will not pair until each orbital contains one electron  according to ​Hund’s rule​, and all such electrons will have the same spin.    2​ 2​ 3​ For example, the electron configuration for nitrogen is 1s​ 2s​ 2p​ . The spins for the electrons in  each orbital are: 1s ⇃↾, 2s ⇃↾, 2p ↾ ↾ ↾.  Because of Hund’s rule, the electrons will occupy each  2p orbital before pairing with another electron, which is why the spins for the 2p orbital are ↾ ↾ ↾  and not ⇃↾ ↾ _.    Chromium and copper are examples of special cases where the 3d orbitals get occupied before  there are two electrons in the 4s orbital, despite the 4s orbital having lower energy.  For  chromium, the electrons in these orbitals look like: 3d ↾ ↾ ↾ ↾ ↾, 4s ↾.  For copper: 3d ⇃↾ ⇃↾ ⇃↾  ⇃↾ ⇃↾, 4s ↾.      Electron Configurations for Ions:    An ​ion​ is an atom with an electric charge as a result of gaining or losing valence electrons.    For example, the electron configuration for a sodium atom is 1s​ 2s​ 2p​ 3s​ .  A sodium atom can  2​ 2​ 6​ 1​ lose an electron and become a sodium ion with a +1 charge.  The lost electron will come from  the valence shell, which means the new configuration will be 1s​ 2s​ 2p​ 3s​  or just 1s​ 2s​ 2p​ . 2​ 2​ 6​ 0​ 2​ 2​ 6​   2​ 2​ 5​ As another example, the electron configuration for a fluorine atom is 1s​ 2s​ 2p​ .  A fluorine atom  can gain an electron and become a fluorine ion with a ­1 charge.  The gained electron will be  2​ 2​ 6​ added to the valence shell, which means the new configuration will be 1s​ 2s​ 2p​ .       


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