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Week 9 Book Notes

by: Cassidy Zirko

Week 9 Book Notes Chem 141

Cassidy Zirko
College Chemistry 1
Mark Cracolice (P)

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College Chemistry 1
Mark Cracolice (P)
Class Notes
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This 11 page Class Notes was uploaded by Cassidy Zirko on Saturday October 31, 2015. The Class Notes belongs to Chem 141 at University of Montana taught by Mark Cracolice (P) in Fall 2015. Since its upload, it has received 34 views. For similar materials see College Chemistry 1 in Chemistry at University of Montana.


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Date Created: 10/31/15
Week 9 Chem 141 Prof Cracolice Chapter 23 How are Electrons Distributed With in an Atom The Bohr Model 102615 231 What is the Nature of Light Electromagnetic Radiation form of energy that consists of both electric and magnetic fields Electromagnetic spectrum gamma rays X rays Ultraviolet light Infrared light microwaves radio waves visible light 300 108 Speed of llght c S Light takes a little more than 8 minutes to travel from the sun to earth Wave equation C 2 I where C speed of light vzfrequency and 2x wavelength The speed of light is a fixed quantity Wavelength and frequency are inversely proportion Longer wavelength I shorter frequency Continuous spectrum band of colors that results from electromagnetic radiation emission over a range of wavelengths Line spectrum spectral lines that appear when light emitted from a sample is analyzed in a spectroscope Discrete lines indicate that elements are individually distinct Each element has a unique line spectrum Relationship between intensity and light emitted I energy of electrons in atoms of a material is directly proportional to frequency E CC V where E energy Photon particle of light Wave particle duality all matter and energy possess both wavelike and particle like properties photons can behave like waves and particles Line spectrum only has light at some wavelengths Wave length is usually expressed in m meters or nm nanometers Solve electromagnetic radiation problems to 3 sig figs 232 What are the Characteristics of Particles of Light Energy of a photon is proportional to its frequency E CC V I E h v where E energy h Planck s Constant Week 9 Chem 141 Prof Cracolice 34 Planck s Constant 6626 10 J 39S specifies that light is not a continous ow of raidation but a series of individual photons Has a meaningful effect on the behavior of energy in small systems Photoelectric Effect light can cause ejection of electrons from some metal surfaces Energy of a photon Energy needed to eject electron kinetic energy of election after ejection Emaic 12 mv2 kinetic energy of an electron Ephoton Eeject Ekinetic l kgm2 2 S l Joule l 31 Mass of electron 910 10 kg Remember to square the speed at which electrons are ejected because then you get mZs2 23 Remember that there are 602 10 electrons in a mole and that Avogadro s number 602 1023 can simple be used as male 233 How are Electrons Arranged in Atom Quantized limited to specific values it may never be between 2 values Continuous can have any values between any two values there is an infinite number of acceptable values Every possessed by election in a hydrogen atom is quantized Line spectrum quantized Spectrum of white light continuous Quantized energy levels at any instant electrons may have one of several possibly energies but at no time can it have an energy between them Electron can orbit at a certain specified distances but never found between those distances Quantum jump quantum leap when an electron moves between orbits Electrons are normally found at ground state Ground state condition when all electrons in an atom occupy the lowest possible energy level 0 I ground state for hydrogen nl energy level Excited State condition and which one or more election in an atom has an energy level above ground state Excited State unstable electron Week 9 Chem 141 Prof Cracolice 0 When electron goes back down to ground sate form excited state photon energy is released proportional to the energy difference between the two levels 0 Energies of electrons increase as they move away from the nucleus 0 Energy of electron is proportional to inverse of sequence of n l E I E 2 RH Rydberg constant for hydrogen RHgtIlt1 n 21798722 10 18 Chapter 24 How are Electrons Distributed within and Atom The Quantum Mechanical Model 102815 241 What are the Short Comings of the Bohr Model of the Atom 0 Some cases light behaves as a wave I when light travels 0 Some cases light behaves as a particle I when produced or interacts with something 0 How Does and Electron Wave Move around the Nucleus I timed mi 00000 Describes the wavelength of any particle Wavelength is a critical component of properties Wave like nature enables us to understand why there are fixed orbits of electrons Orbital circumference is quantized Circumferences have only certain values n1 one wave length per orbit as 11 increases there are greater integers of multiples of the wavelength 0 How Precisely can an Electron in an Atom be Located 0 O Heisenberg I proved that it is impossible to simultaneously know both the position and velocity of an electron There is an uncertainty to each of those two measured quantities h Heisenberg Uncertainty Principle A x l lt m AD 2 E A x uncertianty E position A i uncertianty of velocity h planks constant hand4 H relationship tells us that the product of uncertainty in position multiplied by the uncertain in velocity is greater than a constant smaller uncertainty in one quantity leads to greater uncertainty in the other are constants for a given particles 242 How is the Bohr Model Modified to Account for De Broglie s and Heisenberg s Findings Week 9 Chem 141 Prof Cracolice 0 How Precisely can an Electron in an Atom be located 0 O 0 Quantum Mechanical Model of the Atom an atomic concept that recognizes four quantum numbers by which electrons energy levels may be described Schrodinger s Wave equation used to understand how wavelike properties of electrons effect their behavior Probability Density probability that a particles is found in a specific three dimensional region in space Probability tells the likelihood of finding an electron in a region of space I What is the Role of the Bohr Model in the Quantum Mechanical Model of an Atom 0 O 0 O O Principle energy levels were established nl n2 etc No end to principle energy levels Seventh levels is the highest occupied by the ground state electrons Energy possessed by an electron depends on the energy level Volume of spaces also increase with increasing principle quantum number I How do we Account for Electron Energies Between Principle Energy Levels 00 O OO O 0 specific sublevel identified by principle energy level and sublevel l quantum number for sublevels are designated by values for 12091247 1 mainly l0123 1 corresponds directly to the sublevel total number of sublevelsn nprinciple quantum number I I n1 1 can be only equal to O I one sublevells I n2 1 elements other than hydrogen energy of each principle energy level spreads over can be equal to l or O I two sublevels 2s and 2p a range of sublevels I At n2 increasing order of energy 2slt2p I At n3 3slt3plt3d I At n4 4slt4plt4dlt4f Quantum number 3 and 4 energy ranges overlap I I n3 3d has a higher energy than n4 4s Same sublevel n3 are lower than n4 3slt4s 3plt4p I How does the Quantum Model Describe Electron Locations 0 Not possible to describe path traveled by electron Week 9 Chem 141 Prof Cracolice O Orbitals are essentially probability densities 0 Electron orbital quantum numbers ml mlz l0l I 0 Number of orbitals depends on the electron quantum number I l0 I mz0 I l1 I quotre101 I 2 I meZrlaoalaz l 3 I quotre3 O shape of the s orbital is spherical O principle quantum number increases I size of the orbital increases 0 How Many electrons can occupy an orbital Chapter 25 How are Electrons Distributed Among Orbitals within an Atom 103015 251 How are Electrons Distributed among Orbitals within an Atom 0 In What Sequence do Electrons Fill Orbital s 0 Electron Configuration ground state distribution of electrons among the orbitals of gaseous atoms 0 Atomic number of an atom increase the specific sublevels are filled in differently based on their position on the periodic table O Other sublevels have higher energy but no electrons not occupied 0 Reading periodic table left to right I order of increasing sublevel energy 39 1 period I S sublevel 39 2 d period ZS and 2p 39 3 period 35 and 3p 39 4 period 45 3d and 4p 39 5 period 55 4d and 5p 39 6 period 65 4f 5d and 6p Week 9 Chem 141 Prof Cracolice 00 00 DO 0 0 How can Electron Configuration be Predicted I 7 period 7s 5f 6d and 7 Periodic table is a guide to the order of increasing energy Ground state configurations I Li 132 2S1 I Be ls2 2s2 I Ne ls2 2s2 2p6 First 10 elements electrons will all begin with 1s2 2s2 2p6 Noble Gas Core short hand notation used in electron configuration when the chemical symbol of a noble gas is in square brackets which substitutes for the electron configuration of the element I Ne 1s22s2 2p6 I Electron configuration for sodium Ne 3s1 Works for any and only noble gases Group 8N18 Chromium Ar 4s2 3d5cli1ferent than what it should be because of the stability of half full and completely full orbitals following the normal rules it would be Ar 4s2 3d4 but because an orbital with 4 electrons is very unstable chromium adds a fifth electron to make it stable same goes with elements that would have 9 electrons in the cl orbitals Should be able to reproduce electron ground state configuration based on a table not by memory 0 Number of electrons in highest energy sublevel relates to position of the element of that atom I group 1N1 ns1 where 11 the highest occupied principle energy level 0 Remember the atomic number number of protons number of electrons I How can the Number of Paired and Unpaired Electrons in an Atom be Predicted 0 OOO Orbital diagram diagram that shows how many electron are in each orbital Similar to writing electron configuration Orbital that is occupied by l electron arrow in box pointing upward Orbital that is occupied by 2 electrons 2 arrows in a box one pointing up and one pointing down Week 9 Chem 141 Prof Cracolice Oxygeu Ti H TH T ls 2s 311 0 Make sure to fill each orbital completely before moving on to the next one Then fill each box with one arrow before going back and doubling up on arrows O Hund s Rule the most stable arrangement of electrons is the one that has the maximum number of unpaired electrons 0 What Symbols are used to Represent the Highest Energy Electrons in an Atom 0 Similar chemical properties of elements in the same group are related to the total number of s and p electrons O Valance Electrons the highest energy s and p electrons in an atom which determines the bonding characteristics of an element 0 n principle quantum number ns1 I highest principle energy level for all groups 1A1 elements 0 valance electron number is based on the group number 0 Lewis Dot Symbols electron dot symbols shows the number of valance electrons for an element 0 Element symbol surrounded by correct number of valance electrons O Paired electrons occupy the same orbital usually placed on the same side 0 All 8 electrons fulfill the octet rule only noble gases fulfil this requirement without bonding 39 Cl 39 O O 0 Chapter 26 What Causes Trends in Properties of Elements 11215 261 What Causes Periodic Trends in Properties of Elements 0 Franklard I saw patterns in chemical compounds 0 Valence combining power of an element Week 9 Chem 141 Prof Cracolice Periodic table increasing atomic weight horizontally with the same valance vertically Similar to other correlation of properties What Causes Periodic Trends in Effective Nuclear Charge 0 O O O O 00 Partly dependent on nuclear charge Zeff charge eXperience by an election in an atom with many electrons Negative charges shield some of the positive charge from the nucleus Effective nuclear chare is actual nuclear charge Z atomic number subtracted rom the shielding constant S Zeff 23 Effective nuclear charge can be deduced by qualitatively estimating effective nuclear charge of an element using S shielding effect is looking to see how many possible electrons from inner orbital will shield the outer most electrons EX Li charge Z 3 s outer most electrons in 2s orbital about 2 Zeff 3 21 Its only about 2 because we are unable to know the position of both electrons of the outermost orbital so we can conclude that the 2 outer most electrons wont always be shielded from the nucleus Electrons in the same principle energy level do not shield as well as those in lower principle energy levels 211d period elements Li Be B C N O F Ne Approximate Zeff l 2 2 3 4 4 5 6 0 Effective nuclear charge I increases across periodic table 0 For group lAl S is equal to the noble gas core for the number of electrons minus 1 because ns1 electrons spend some of their time closer to the nucleus then all of the lower energy levels What Causes Periodic Trends in the Size of Atoms 0 Atomic Radius average distance between the nucleus of an atom of the element and the outer limits of the electron cloud Atomic orbital probability of finding an electron not a definite boundary Heisenberg uncertainty principle limits the ability to measure the distance between nucleus and the outer most electron of an atom Atomic radius increases as you go down a group and decreases as you go across from left to right a period Highest occupied energy level and effective nuclear charge shows reasoning for trend in atomic size Summary Atomic Size I Increases from right to left across any row of the periodic table and from top to bottom smallest atom are towards upper right corner and the larger are toward the bottom left corner What Causes Periodic Trends in the Size of Ions Week 9 Chem 141 Prof Cracolice Trend among metal elements are that cations of an element are smaller than the original atom Element radius is bigger then ion of the element radius Reduced electron electron repulsion with ion I same positive charge but exerted on fewer electrons Anion is larger than the original atom because there increase electron electron repulsion Size of atom and size of radii increases in the same way Isoelectric series group of atoms andor ions that have the same number of electrons iso equal Summary Ionic Size I Size of ions increases down a group in periodic trend size of ion in an isoelectric series decreases with the increasing atomic number Monoatomic cations are smaller than original atom Monatomic anion I larger than the original atom What Causes Periodic Trends in Ionization Energy O 0000 CO Ionization energy energy required to remove one electron from a neutral gaseous atom of an element Atomic number increases I ionization energy increases with in a period Ionization energies are lower as the atomic number increase within a group Increase going left to right across the table Higher Zeff I negatively charged valance electrons are attracted more strongly to the nucleus I more energy to remove electrons Second ionization energy energy needed to remove second electron Third ionization energy energy needed to remove third electron Summary First Ionization energy I Generally increases from left to right and from the bottom to the top of the period table The highest first ionization energy is in the upper right hand corner What Causes Periodic Trends in Electron Affinity O 0000 Electron Affinity change in energy that occurs when an electron is added to a gaseous atom or ion Neutral atom accepts electrons and from anion I exothermic A E Greater electron affinity I more energy released No specific trend Summary Electron Af nity I General affinity increases from right to left on any row atoms with most electron affinities are on the right side I Increasing energy released I bigger number to smaller I Affinity numerically smaller number to larger number What Causes Trends in Metallic Character 0 Metal can lose one or more electrons and become positively charged Week 9 Chem 141 Prof Cracolice O Nonmetal lacks the metal quality 0 Metallic character increases as you go down a column 0 Smaller atom I harder to let go of electrons O Metalloids semimetals properties of both metals and nonmetals O Stair step line between carbon and down to polonium separates metals and nonmetals elements along this line are semimetals 0 Summary Metallic Character 39 Increases from right to left across the periodic table Metals Nonmetals Loses electrons easily to form cations Tend to gain electrons to form anions 12 or 3 valance electrons 4 or more valance electrons Low ionization energy High ionization energy Forms compounds with nonmetals but not Forms compounds with metals and nonmetals with metals High electrical conductivity Poor electrical conductivity High thermal energy Poor thermal energy Malleable Brittle Ductile Nonductile 0 What do Some Groups of Elements from chemical Families 0 Chemical Families groups with properties in common 0 Main groups of elements 0 Alkali metals 0 Alkali Earth Metals O Halogens Valance Electrons ns1 Group lAl minus hydrogen Easily lose valance electrons I similar properties Ionization energies decrease as atomic number increases Reactivity tendency to react with other elements to form compounds Don t normally look like metals Can combine with oxygen to oxidize Density increase going down the table Valance electrons ns2 First and second ionization energy are low Trends are similar to alkali metals Reactivity increases going down table Group 2A2 Less physical property trends Valance electrons nsznp5 Group 7Al7 Also known as salt formers 7 valance electrons Very easily gain an electron Reactivity decreases down the table uorine is the most reactive Density melting and boiling point increase going down the table Week 9 Chem 141 Prof Cracolice O Noble Gases Valance electrons nsznp6 Very unreactive High ionization energies have full outer orbital 0 Hydrogen Valance electrons ns1 Not an alkali metal or halogen Gains electrons to form ions with a positive charge I Has properties of both alkali metals and halogens


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