Principles of Chemistry I
Principles of Chemistry I CHEM 1307
Popular in Course
verified elite notetaker
Popular in Chemistry
This 36 page Class Notes was uploaded by Kara Dibbert on Thursday October 22, 2015. The Class Notes belongs to CHEM 1307 at Texas Tech University taught by Tamara Hanna in Fall. Since its upload, it has received 52 views. For similar materials see /class/226511/chem-1307-texas-tech-university in Chemistry at Texas Tech University.
Reviews for Principles of Chemistry I
Report this Material
What is Karma?
Karma is the currency of StudySoup.
Date Created: 10/22/15
The Structure of Atoms Electromagnetic Radiation Radiation that consists of wavelike electrons and magnetic fields including light microwaves and Xrays Wavelength 7L The distance between successive crests or troughs of a wave The units of length are typically meters or nanometers Frequency v The number of cycles the wave undergoes in one second The unit is typically s1 or Us and is called a hertz Hz Electromagnetic Radiation Wavelength and frequency are related to the speed c at which the wave is propagated units of ms speed of light 30 x 108 ms The electromagnetic spectrum shows all types of different electromagnetic radiation related by A and v Energy Increases As v increases 9 decreases As 9 increases v decreases 10 10 1020 10 10 6 10 10 10 quotJ 10quot 106 10quot 102 10 u Hz l 1 l l l l l l l I l J y rays xrays UV IR Microwave FM I AM Long radio waves Radiowaves I l I I I I l I I 10 1quot 1o 1 1o 2 10 1 10 3 To 10 4 10 2 10 102 10quot 106 108 A m Visible spectrum I Lon 7 low E I Short 2 hI h E 400 500 600 700 A nm Energy increases Wavelength increases BrooksICoIe Cengage Learning Electromagnetic Radiation Some diamonds appear yellow because they absorb purple light of frequency 723 x 1014 Hz Calculate the wavelength of the absorbed light in units of nm 10399 m and A 103910 m C7w C 7 v c 30 x 108 mm 397 A V 723X1014fl 415x10 m 1 nm 415X 10397 m X 415X 102 nm 415 nm 10399m 1A 415x107m x 415x103A 103910 m The Ultraviolet Catastrophe When a metal is heated EM radiation is emitted in wavelengths that depend on the temperature As T increases the red color brightens and eventually white light is emitted Our eyes detect the radiation that occurs in the visible region of the EM spectrum but other wavelengths are also given off by the heated object At the end of the 19th century scientists could not explain the relationship between intensity and wavelength for radiation given off by a heated object They predicted the intensity should 104 increase continuously with decreasing wavelength never reaching a maxima 1 K The theories failed In the ultraVIolet 1000 2000 3000 region short A Wavelength 7 nm Quantization and Planck s Constant Max Planck described that emitted EM radiation was caused by vibrating atoms oscillators and each oscillator had a fundamental frequency of oscillation Energy is quantized such that only certain energies are allowed Planck s constant h 6626 X 103934 Js If the oscillator changes from a higher energy to a lower one energy is emitted as EM radiation AE Anhy If An 1 ilower E level 200 300 400 500 600 700 500 900 1000 wavelengthum Wavelength n rn Intensity of Emitted Light gt Intensity of Emitted Light gt o Einstein and the Photoelectric Effect Electrons are ejected when light strikes the surface of a metal if the frequency of the light is high enough threshold freq If the frequency is too low no electrons are ejected If the frequency is high enough increasing the intensity of the light causes more e39 to be ejected bc more photons of light Light has particlelike properties called photons which are packets of energy Photons are massless and have E proportional to the frequency of radiation Metal only loses electrons if photons have enough energy VlV Using Planck s Equation CD players use lasers that emit red light with a wavelength of 685 nm What is the energy of one photon of this light What is the energy of 1 mol of photons of red light Convert 7K to v c 30X108I7 s 14 1 v T 685X109W 438X10 5 Calculate the E per photon E hv 6626 x 1034 JAphoton438 x 1014 3 1 290 x 103919 Jphoton Calculate the E per mole E 290 x 1019 JW602 x 1023 Wsmol 175 X 105 Jmol 175 kJmol Atomic Line Spectra If a high voltage is applied to atoms of an element in the gas phase at low pressure the atoms absorb energy and become excited The excited atoms can emit light The light from an excited atom consists of only a few different wavelengths of light The line emission spectrum of each element is different Gas discharge tube a i g I quotta quot5 hydro e Atomic Line Spectra Every element has a unique spectrum which can be used to identify an element or determine how much is present 500 I Hg I NE Why do excited gaseous atoms emit light of only certain frequencies Balmer Equation John Balmer found an equation that could calculate the wavelength of red green and blue lines in the visible emission spectrum of hydrogen For hydrogen n1 gt 2 Rydberg constant R 10974 x 107 m391 UV series n1 1 n2 2 3 4 Vis series n1 2 h2 3 4 IR series n1 3 n2 4 5 6 Ultraviolet Visible infrared series series series 395 r A 1 f A O 200 400 600 800 1000 1200 1400 1600 1800 2000 nrr Bohr Model of the Atom Classic physics in the early 20th century described that electrons moving in the positive electric field of the nucleus eventually lose energy and crash into the nucleus If this was true matter would selfdestruct Niels Bohr proposed a planetary structure for the atom where electrons move in a circular orbit around the nucleus much like a planet revolves around the sun Bohr Model of the Atom Bohr said that electrons move around the nucleus only in certain allowed circular orbits Each orbit is associated with a definite energy level Only certain energy levels are allowed quantization The energy possessed by a single electron in the nth orbit energy level of an atom can be described by En energy of the electron Jatom Rydberg constant R 10974 x 107 m391 Planck s constant h 6626 X 103934 Js Speed of light c 30 x 108 ms Principal Quantum Number n Electron Energies Calculate the energy for the transition of an electron from quot5 39 the n 1 level to the n 4 level of a hydrogen atom quot3 Calculate the E when n 1 n1 E Rhc 1097 x 107 ryl16626 x 1034 Jz 30 x 108 my 1 2 n2 1 218 x 103918 Jatom Calculate the E when n 4 Rhc 1097 x 107 ryl16626 x 1034 Jz 30 x 108 a A 4 2 2 n E is absorbed when an e39 goes 36 X 1019 Jatom from a lower to higher E level Calculate AE for the process H AE E4 E1 136 x 1019 218 x 1018 204 x 1018J Change in Energy Levels If an electron moves from one energy level to another energy must be either absorbed or released If an electron is excited from n 1 to n 2 AE Ef Ei NARhcZZ NARhc12 075NARhc 984 kJmol NA Avogadro s Number Therefore moving an electron from the first to the second energy state requires an input of 984 kJmol of atoms The electron can only be excited at this precise amount of energy QUANTIZATION Moving an electron from a higher energy state to lower energy state leads to the release or emission of energy ien2 gt n 1AE984kJmol Change in Energy Levels Qualitative Classify each of the hydrogen transitions n r nz n x n x A 11 4 T 14 14 a 174 113 115 L 113 13 111 nzl 111 111 j A B c D Absorption AC Electron goes from a lower to higher energy level Emission B D Electron goes from a higher to lower energy level Ionization C Electron goes from the lowest energy level to infinity The electron is moved completely away from the nucleus Bohr Model of Hydrogen Electrons are excited to higher energy levels and absorb energy The electrons can then return to any lower energy level either directly or in a series of steps and release energy as photons of EM radiation The change in energy for the emission lines in excited hydrogen atoms can be calculated as follows Electronic Transitions of Hydrogen n Energy Jatam Lyman series Balmer series Lyman Series UV region 4 e39movesfromn gt 1ton 15 z Balmer Series Visible region e39movesfromn gt 2ton 22 IR region 6 moves from higher levels to n 3 4 or 5 a OOOOOOOO aaaaaaaaaa m Invisible Lines Ultraviolet 410 2 rim 434 1 nm 4861 Hm 556 3 rim Invis b lines Infrared Particle Wave Duality In 1925 de Broglie proposed that electrons behave as both particles and waves A moving particle has an associated wavelength An electron can behave like a wave or a particle but no experiment shows a wave and a particle simultaneously Calculate the wavelength associated with an electron of mass m 9109 X 103931 kg that travels at 40 the speed of light 1 6626 X 103934 kgmZs mv 39 9109 x 1031 kg12 x 108 ms A 607 X 103912 m Heisenberg s Uncertainty Principle For an electron in an atom it is impossible to determine accurately both its position and energy If we know the energy of the electron with a small uncertainty there is a large uncertainty on its position Chemists predict the approximate location of an electron Schrodinger s Wave Function Schrodinger developed a model for electrons in atoms using the idea that electrons could behave as waves The model uses mathematical equations of wave motion to generate wave functions LP Each describes an allowed energy state of an electron Schrodinger s Wave Function An electron can be described as a standing wave Only certain vibrations are possible not2 Vibrations are quantized where n is a quantum number Nodes occur at points of zero amplitude 12 a BreaksCull Gen aaaaaaaaa rig Schrodinger s Wave Function An electron in 3D space requires three quantum numbers n e and me all are integers Each LP is associated with an allowed energy level L1 2 is proportional to the probability of nding an electron at a given point Orbitals describe the region of space where an electron of a given energy is most likely to be located 121s Distance r from nucleus Quantum Numbers Quantum numbers n are a set of positive numbers with integer values that define the E of the allowed orbits in an atom The lower the value of n the smaller the orbit the closer the electron is to the nucleus the lower the energy state The energy of an electron in an orbit has a negative value Ground state The state of an atom in which all the electrons are in the lowest possible energy levels For hydrogen n 1 Excited state The state of an atom in which at least one electron is not in the lowest possible energy level For hydrogen n gt 1 E exchange occurs when electrons move to different orbit levels Quantum Numbers Any atomic orbital is specified by three quantum numbers n Principle Quantum Number n 1 2 3 Size of the orbital distance from the nucleus The higher the n the higher the energy level Electrons with the same n are in the same electron shell 6 Azimuthal Quantum Number 6 0 1 2 n1 Shape of the orbital Value is limited by n When multiple 6 values are possible called subshells Value ofe Subshell Label 0 s 1 p 2 D 3 f Quantum Numbers Any atomic orbital is specified by three quantum numbers me Magnetic Quantum Number me 0 I1 I2 Ie Orientation in space of orbitals in a subshell 6 determines me Number of orbitals in a subshell 26 1 Value ofe Subshell Label Value of me 0 s 0 1 p 1 0 1 2 cl 2 1 0 1 2 3 f 3 2 1 0 1 2 3 Shells and Subshells n2 number of orbitals in the shell The First Electron Shell n 1 n 1 e 0 m 0 one s orbital denoted ls In the shell closest to the nucleus only a single orbital exists The Second Electron Shell n 2 n 2 e 0 m 0 one s orbital denoted Zs n 2 e 1 m 1 0 1 three p orbitals denoted 2p The p orbitals have the same shape but different orientation The Third Electron Shell n 3 n 3 e 0 m 0 one s orbital denoted 3s n 3 e 1 m 1 0 1 three p orbitals denoted 3p n 3 e 2 m 2 1 0 1 2 five d orbitals denoted 3d Relationship Among Quantum Numbers of Total of Possible Subshell orbitals in orbitals in n value ofe Designation Possible value of m subshell a shell 0 ls 0 1 1 2 0 ZS 0 1 1 2p 1 0 1 3 4 3 0 3s 0 1 1 3p 1 O 1 3 2 3d 2 1 O 1 2 5 9 4 0 4s 0 1 1 4p 1 0 1 3 2 4d 2 1 0 1 2 5 3 4f 3 2 1 0 1 2 3 7 16 Quantum Number Practice What type of orbital is described by the quantum numbers Q n3e2me2 A 3d orbital Qn5e0me0 A 5s orbital What shells contains 4 orbitals n 2 n2 orbitals in a shell In what shell do f orbitals begin n 4 The Shape of Atomic Orbitals No more than 2 electrons can be held in an orbital s orbitals e 0 Spherical in shape As n increases the size of the spherical orbital increases so the electrons move farther away from the nucleus An electron cloud picture shows the probability of finding an e39 at a distance r from the nucleus The e39 will never be at the nucleus 90 probability of finding the e39 inside this sphere Most probable distance of H 15 electron from the nucleus 00529l nm Probability of nding electron at given distance from the nucleus 9 1 2 3 4 5 6 Distance from nucleus 1 unit 00529 nm The Shape of Atomic Orbitals No more than 2 electrons can be held in an orbital p orbitals e 1 Shape resembles a dumbbell e39 spends time in both lobes A nodal surface passes through the nucleus no probability of finding an electron Three p orbitals are in a subshell all have the same shape but lie along a different axis px py p2 Radial probability distribution px py 22 r 10 m Brooksr Cole Cengage Leamang The Shape of Atomic Orbitals No more than 2 electrons can be held in an orbital d orbitals e 2 Four orbitals shaped like cloverleafs one shaped differently For the cloverleaf orbitals two nodal surfaces pass through the nucleus so there are four regions of e39 density The orbitals lie along two planes The Shape of Atomic Orbitals No more than 2 electrons can be held in an orbital f orbitals e 3 Seven total orbitals Three nodal planes Electron density lies in eight regions of space quot R 4fxz 7 43y2 2 The Periodic Table Can Help il s M k eliemen ts d lhllnc u elements trans unu metals I Pb lmke1lement5 f b uck e ements Eanfthan des and a c n ides Sf Electron Spin ms Electrons can be thought of as having a spin which allows them to be attracted or repelled by a magnetic field If an unpaired electron is placed in a magnetic field there are two orientations for the atom 1 when the electron spin is aligned with the field and 2 when the spin is opposed to the field ms Electron Spin Quantum Number n1S 12 12 Electron spin is quantized The electron can either be align or opposed to a field L M Hi i i H 0 N O Tig i unpaired electron He 0 Filledshell inertgas O o 1s 1s 2 2p H2 H H H H l L39 9 F Osaka mama Be 0 gt paired electrons 9 O 7 g g atten inert gas H H r Hillil llr3E B trait 753 l l l l l i l l C MQOEEE p g Magnetism Dia magnetism All electrons in an element or compound are paired When placed in a magnetic field the substance experiences slight repulsion Paramagnetism At least one electron in an element or compound is unpaired In the absence of a magnetic field the electrons are randomly oriented When placed in a magnetic field the substance experiences attraction 39 39 Magnetism Ferromagnetism A form of paramagnetism where the magnetic effect is greatly enhanced In the absence of a magnetic field the electrons align themselves in the same direction When placed in a magnetic field the substance experiences a strong attraction Very few metals such as iron cobalt and nickel exhibit these properties