Chemistry Chapter 3 Notes
Chemistry Chapter 3 Notes CHEM 1030 - 002
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This 6 page Class Notes was uploaded by Robert Kessinger on Monday September 14, 2015. The Class Notes belongs to CHEM 1030 - 002 at Auburn University taught by Walter Felix Casper in Summer 2015. Since its upload, it has received 88 views. For similar materials see Fundamentals Chemistry I in Chemistry at Auburn University.
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Date Created: 09/14/15
Chapter 3 Notes Energy and Energy Changes 0 Energy is the capacity to do work or transfer heat Kinetic Energy results from motion m is mass u is velocity El2muA2 A type of kinetic energy is thermal energy associated with random motion of atomsmolecules Can be monitored through temperature changes 0 Potential energy is possessed by an object by virtue of its position Chemical energy is stored within the structural units of chemical substances Electrostatic energy results from the interaction of charged particles EQlQ2d 0 Kinetic and potential energy are interconvertible 0 Although energy can assume many forms the total energy of the universe is constant Energy can neither be created nor destroyed When energy of one form disappears the same amount of energy reappears in another form or forms 0 This is known as the law of conservation of energy The Unit of Energy 0 The SI unit of energy is the joule J named for the English physicist James Joule 0 It is the amount of energy possessed by a 2kg mass moving at a speed of 1 ms E 12 mu2 122 kgl ms2 1 kgm2s2 1 J The joule can also be defined as the amount of energy exerted when a force of l newton N is applied over a distance of 1 meter 1 J l N m Because the magnitude of a joule is so small we often express large amounts of energy using the unit kilojoule kJ l k 1000 J Properties of Waves Wavelength A lambda the distance between identical points on successive waves Frequency 1 nu the number of waves that pass through a particular point in 1 second Amplitude the vertical distance from the midline of a wave to the top of the peak or the bottom of the trough The speed of light 6 through a vacuum is a constant c 299792458gtlt108 ms 0 Normally rounded to c 300gtlt108 ms Speed of light frequency and wavelength are related CAV 0 A is expressed in meters 12 is expressed in reciprocal seconds s39l s391 is also known as hertz Hz The Electromagnetic Spectrum From most energy to least Gamma Xray UV purple light blue light green light yellow light orange light red light Infrared Microwave and radio waves Quantum Theory Early attempts by nineteenthcentury physicists to gure out the structure of the atom met with only limited success They were using the laws of classical physics These laws describe the behavior of macroscopic objects Over time the realization and acceptance was the behavior of subatomic particles is not governed by the same physical laws as larger objects Quantization of Energy When a solid is heated it emits electromagnetic radiation known as blackbody radiation over a wide range of wavelengths The amount of energy given off at a certain temperature depends on the wavelength Ex coals in a fire Classical physics assumed that radiant energy was continuous that is could be emitted or absorbed in any amount Max Planck suggested that radiant energy is only emitted or absorbed in discrete quantities like small packages or bundles A quantum of energy is the smallest quantity of energy that can be emitted or absorbed The energy E of a single quantum of energy is Ehv his called Planck s constant 663x10 34 J s The idea that energy is quantized rather than continuous is like walking up a staircase or playing the piano You cannot step or play anywhere continuous you can only step on a stair or play on a key quantized Photons and Photoelectric Effect 0 The photoelectric effect states that electrons are ejected from the surface of a metal when exposed to light with a threshold frequency The number of electrons ejected is proportional to the intensity brightness of light Below the threshold no electrons are ejected Einstein Beam of light is actually a stream of particles called photons Ephotonhv hvKE W White Light and the Emission Spectra Sunlight is actually composed of various colors energies Energizing a material will let it show it s emission spectra Atomic Line in Spectra 0 Line spectra is the emission of light at only certain wavelengths Every element has a line spectra unique to it 1 AR1n1 quot21n2quot2 Bohr s Model Electrons in an atom can only occupy certain orbitals n they are quantized Electrons in permitted orbits have specific allowed energies Energy is absorbed or emitted only as the electron moves from one allowed energy state to another High to low energy emits a photon Bohr showed that the energies of the electron in a hydrogen atom are given by the equation Energy of allowed orbit En218 x 10A18 J 1nquot2 En is the energy n is a positive integer As an electron gets closer to the nucleus n decreases En becomes larger in absolute value but more negative as n gets smaller 0 E is most negative when n 1 Called the ground state the lowest energy state of the atom For hydrogen this is the most stable state 0 The stability of the electron decreases as 11 increases Each energy state in which n gt 1 is called an excited state The Line Spectra of Hydrogen 0 11 is the initial state 0 m is the final state 1 L AE 1112 218 gtlt 1018 J 2 it mi N 39 h A photon is emitted when ni gt m AE is negative energy is lost to the surroundings 0 A photon is absorbed when m gt ni AE is positive The energy difference between the initial and final states is l 218 gtlt1018J l 1 2L he n n Wave Properties of Matter Bohr s theory fit experiment but why was the electron restricted to orbiting the nucleus at fixed distances Louis de Broglie reasoned that if light can behave like a stream of particles photons then electrons could exhibit wavelike properties standing waves The de Broglie Hypothesis Only certain wavelengths allowed with a node at zero 11 mu 0 A is the wavelength associated with the particle m is the mass in kg u is the velocity in ms Quantum Mechanics Waves could have matter like properties and matter could have wave like properties 0 The distinction was no longer clear wave particle duality Heisenberg Uncertainty Principle 0 It is impossible to know simultaneously both the momentum and position of a particle The more precise one measurement is the less the other becomes Axis the uncertainty in position in m 0 Ap is the uncertainty in momentum AX ApE h 4T Au is the uncertainty in velocity in ms m is the mass in kg Ax mAu E h 4n The Schrodinger Equation 0 Erwin Schrodinger derived a complex mathematical formula to incorporate the wave and particle characteristics of electrons Wave behaviour is described With the wave function 11 0 The probability of finding an electron in a certain area of space is proportional to 112 and is called electron density Quantum Mechanics defines the region Where the electron is most likely to be at a given t1me The Schrodinger Equation and The Quantum Mechanical Description of the Hydrogen Atom The Schrodinger equation specifies possible energy states an electron can occupy in a hydrogen atom The energy states and wave functions are characterized by a set of quantum numbers Instead of referring to orbits as in the Bohr model quantum numbers and wave functions describe atomic orbitals Electrons are located in orbitals the distribution of the electron density is the probability of locating the electron there Quantum Numbers Describe the distribution of electron density in an atom Derived from the solution to Schrodinger Equation for the hydrogen atom The principal quantum number n designates size The angular moment quantum number 1 describes shape The magnetic quantum number m1 specifies orientation Principle Quantum Numbern The principal quantum number n designates the size of the orbital 0 Larger values of n correspond to larger orbitals The allowed values of n are integral numbers 1 2 3 and so forth 0 The value of n corresponds to the value of n in Bohr s model of the hydrogen atom A collection of orbitals With the same value of n is frequently called a shell Angular Momentum Quantum Numberl The angular moment quantum number 1 describes the shape of the orbital The values of l are integers that depend on the value of the principal quantum number The allowed values of 1 range from 0 to n l gt Example Ifn 2 1 can be 0 or 1 A collection of orbitals With the same value of n and l is referred to as a subshell Magnetic Quantum Numberml The magnetic quantum number m1 describes the orientation of the orbital in space The values of m are integers that depend on the value of the angular moment quantum number l0l For a value of 1 there are m1 211 values of m The number of 1111 values indicates the number of orbitals in a subshell With a particular 1 value each value refers to a different orbital Electron Spin Quantum Numberms The electron spin quantum number ms is used to specify an electron s spin There are two possible directions of spin Allowed values of ms are 12 and 12 Summary of Quantum Numbers To summarize quantum numbers principal n size angular 1 shape magnetic m1 orientation Atomic Orbitals Technically orbitals don t have welldefined shapes since the wave function characterizing the orbital extends from nucleus to infinity s Orbitals spherica1 in shape but different in size P 2 spherical together 1 2 together With one in middle or clover shape Electron Con gurations 0 Hydrogen has only 1 electron s orbital 0 In many electron systems knowing the electron con guration is needed to describe how the electrons are distributed