Study Guide for Exam 4
Study Guide for Exam 4 CHEM 120
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This 0 page Study Guide was uploaded by Leslie Pike on Friday November 27, 2015. The Study Guide belongs to CHEM 120 at Western Kentucky University taught by Dr. Darwin Dahl in Summer 2015. Since its upload, it has received 88 views. For similar materials see College Chemistry I in Chemistry at Western Kentucky University.
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Date Created: 11/27/15
Study Guide for Exam 4 to be administered December 2 2015 Covers Ch 7 8 9 Chapter 7 Quantum Theory and Wavelengths Very important equation for calculating frequency and wavelength of light waves Speed 2 Frequency Wavelength For all light waves speed is a constant 3108 ms Given speed and frequency one can calculate wavelength or given speed and wavelength one can calculate frequency Frequency is the number of wavelengths of that wave that pass a xed point in one unit of time Frequency is represented by the Greek letter nu which looks like the English letter v A high frequency means that more wavelengths pass the xed point per time interval High frequency is equivalent to high energy The energy of a light wave is equal to Planck39s constant 663103934js multiplied by frequency The unit of frequency is Hz or Us If something has a frequency of 100000 hertz that means that 100000 wavelengths pass through a point in one second Wavelength is selfexplanatory the length of a wave from crest to crest or trough to trough Wavelength is represented by the Greek letter lambda Long wavelengths correspond to low energy and short wavelengths correspond to high energy Wavelength is often measured in angstroms One angstrom is 103910 meters Niels Bohr came up with the energy shell model for the atom This is NOT the Rutherford planetary model and the electrons do NOT quotorbitquot the nucleus like planets orbiting the sun Electrons oat around in their set energy levels Light is emitted when an electron jumps from a higher energy level to a lower energy level Light is absorbed when an electron jumps from a lower energy level to a higher energy level The energy of an electron is equal to the negative Rydberg constant Dr Dahl will provide this on exams divided by the quantum number squared The change in energy when an electron changes states is described by the following equation 1 1 AERH 2 2 I ll 11 Note THIS IS BACKWARDS from the way these normally work In this case we have INITIAL MINUS FINAL instead of the other way around The Balmer series is when an electron jumps from a higher energy state to n2 Light in the visible spectrum is emitted during this jump The Schrodinger wave equation describes the probability of where an electron will be at any given time This equation itself is beyond the scope of a 100level class however you will be expected to know its four variables n l m and ms Nprincipa quantum number It can be any positive integer It represents the energy level or the shell o L angular momentum quantum number It can be any whole number up to N1 It represents whether the subshell is s p d or f The s subshell has an angular momentum quantum number of 0 p is 1 d is 2 f is 3 IMPORTANT THING TO WATCH OUT FOR the s subshell is ZERO not ONE p is 1 etc Dr Dahl will probably try to trip you up on this Mmagnetic quantum number It ranges from to 0 to it can be any integer value between these values It represents the orbital for example px py and pz 0 MS is the spin It can be 12 or 12 Chapter 8 Electron con gurations The ground state of an atom means that the atom is not an ion it has an equal number of protons and electrons The ground state also means that there are no excited electrons The ground state electron con guration for magnesium is 1522522p63sz This can be abbreviated as Ne3sz because ground state magnesium has a neon core and then an extra 35 subshell When subshells ll one electron is put in each orbital rst Then a second electron is put in each orbital Subshells ll in order of energy level The n1d subshell lls before the hp subshell 4s lls then 3d lls then 4p lls There are exceptions to this rule namely when the d subshell has 4 electrons and when the d subshell has 9 electrons In these cases one of the electrons in the s subshell will go back into that last d orbital When subshells empty they empty from the highest nvalue rst 45 empties BEFORE 3d empties NOTE THE ORDER IN WHICH SUBSHELLS EMPTY DOES NOT NECESSARILY MATCH THE ORDER IN WHICH SUBSHELLS FILL If nvaIues are the same the atom Ioses from the highest vaue rst 3d disappears before 3p and 3s A paramagnetic atom has at least 1 unpaired electron A diamagnetic atom has no unpaired electrons Noble gases are diamagnetic as are groundstate alkali earth metals A good way to remember this is to remember that quotdiquot means two and in a quotdiamagneticquot atom each orbital has two electrons in it All elements want to achieve noble gas con guration This means that they have all of their shells full When alkali metals exist in the 1 state they have the same con guration as the noble gas right behind them Na has the same electron con guration as Ne Some atoms achieve pseudo noble gas con guration instead this means that they have the noble gas core and also a lled nd subshell When halogens exist in the 1 state they have the same electron con guration as the noble gas in front of them F39 has the same electron con guration as Ne Elements with the same electron con guration are known as isoeIectronic Ne Na F39 and 0239 are all isoeIectronic The main group aka representative group consists of the sgroup and the pgroup ie IA IIA IIIA IVA VA VIA VIIA and VIIIA The main group does NOT contain the d block or the lanthanides or the actinides Nonmaingroup elements may not necessarily obey the following periodic table trends Periodic tabIe trends and why 0 Atomic radius increases from top to bottom and right to left Reason a new shell is added each period and each shell is bigger than the previous sheII Also as atomic nuclei increase in charge they draw the electrons in closer causing sheII size to shrink Metallic trend increases from top to bottom and right to left This is an easy one to remember the periodic table is labeled in terms of metals metaIIoids and nonmetaIs Electron af nity or the amount of energy required to remove an added eIectron increases from bottom to top and left to right Reason the fewer n IeveIs an atom has the closer the valence shell is to the nucleus and the more strongly the nucleus attracts those valence electrons Also the closer the shell is to full the more strongly the atom feels its missing electrons and the worse it wants them Electronegativity is how much that atom hogs the shared electrons when it bonds with another atom Same trend as electron af nity and for the same reasons NOTE The electronegativity trend for the rst two periods of the p block is that the top period has higher electronegativity than the bottom period when the diagonal is drawn from left to right top to bottom It is the other way around the rest of the time 0 Reactivity This is a weird trend For alkali metals and alkali earth metals the reactivity increases down the column This is because as the valence she gets farther from the nucleus it is less attracted by the nucleus and those outer shell electrons are more easily Iost For hangens and chaIcogens the uorine and oxygen columns reactivity decreases down the column This is because as the valence she gets farther from the nucleus it is less attracted by the nucleus and the atom has less of a tendency to steal other atoms39 eIectrons Chapter 9 Ionic and Covalent bonding If two atoms participating in a bond have an electronegativity difference of 0 the bond is nonpoIar covaIent The atoms share the electrons evenIy If the electronegativity difference is between 0 and 17 the bond is polar covalent they technicaIIy share eIectrons but the more electronegative atom hogs the electrons most of the time If the electronegativity difference is greater than 17 the bond is ionic meaning that the more electronegative element completely steals the electrons The lattice energy is the energy required to completely separate one mole of solid ionic compound into gaseous ions Smaller atoms have a greater Iattice energy Ions with a larger charge have a higher lattice energy A Lewis structure is a diagram of bonds and electrons in a molecule Molecules want to be symmetrical as much as possible keep this in mind when drawing a Lewis structure Any two bonding atoms can form single bonds Carbon oxygen nitrogen phosphorous and sulfur can form double bonds Carbon and nitrogen can form triple bonds As the bond order increases secondorder bond is a double bond and thirdorder bond is a triple bond the length of the bonds decreases Polar covalent bonds are symbolized by an arrow dipole The arrow points toward the more electronegative element Polar covalent bonds do not necessarily mean that the molecule is polar often the dipoles cancel each other out Resonance is what happens when more than one equivalent Lewis structure can be drawn for a molecule Examples benzene ozone carbonate In these cases the atoms quotsharequot the extra bonds The symbol for resonance is DD an arrow with two ends The formal charge of an atom is calculated by counting a lone pair for 2 and a bonding pair for 1 For an atom with a formal charge of 0 the charge equals the charge from the nucleus When multiple Lewis structures can be drawn for a molecule the correct Lewis structure is the one in which the atoms have the smallest formal charges When multiple Lewis structures with the same formal charges can be drawn for a molecule the most electronegative atom should have the negative charge or the most electropositive atom should have the positive charge etc Exceptions to the octet rule Hydrogen and helium only have two electrons in their outermost shell Alkali and alkali earth metals do not have enough electrons in their valence shell to form enough bonds to ll an octet Instead they form bonds with the valence electrons that they have and they do not have full octets This same situation can happen for MA elements such as boron Molecules with a principle quantum number greater than 2 can form extra bonds on account of the dorbital Example sulfur hexa uoride Molecules involving noble gases typically break the octet rule on account of the noble gas already having a full octet The noble gas will have extra electron pairs If the noble gas has a dsubshell the extra pairs are not a problem Argon krypton xenon and radon can form bonds Helium and neon CANNOT form bonds because they do not have a dshell Molecules with odd numbers of electrons break the octet rule Example nitrogen monoxide Nitrogen and oxygen are doublebonded together oxygen has two pairs of nonbonding electrons nitrogen has a nonbonding pair and a lone electron The delta H for a reaction can be calculated by subtracting the energy of all of the bonds of the products from the energy of all of the bonds of the reactants
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