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CHEM 112

Unit 1-Atoms Electromagnetic Radiation -Electric and magnetic fields  perpendicular and in phase -Waves are characterized by wavelength (distance between two crests) frequency (number of oscillations per second) inverse to wavelength speed  same speed in a vacuum -Interference of waves when amplitudes of two waves add together -waves exactly in phase  constructive interference  amplitude increases -waves exactly out of phase  peak of one wave aligns with valley of other wave  amplitudes cancel and become zero -intermediate degrees of interference may occur Visible spectrum (between UV and infra-red): 390 nm (violet) – 790 nm (red) -gamma rays  x rays  UV  infrared  microwave  radio/TV All elements have their own distinctive atomic spectra Photoelectric Effect: emission of electrons when frequency of the incident light (EM) exceeds a specific threshold frequency -Stopping voltage is at 0 until threshold frequency is reached -Higher intensity of EM waves= more ejected electrons but not higher energies Quantum explanation of photoelectric -Photon hits surface, each photon carries E=hv of energy, 1 photon collides with 1 electron and transfers E=hv to the electron -if E=hv exceeds binding energy of electron, then electron is emitted -leftover energy becomes KE of electron Quantum Theory -Energy absorbed/emitted in units called quanta. Energy of 1 quantum is E=hv Allowing only discrete changes in energy: quantization Quantization affects probability a wave will be emitted -lower frequency=lower energy -lower energy=more likely for wave to emit -Absorption and emission causes system to jump between energy levels -input: energy increases by nhv where n is a positive integer (addition of heat) -output: energy decreases by nhv where n is a positive integer, emitted wave has energy E=nhv Quantized model: finite amount of energy in the box, yields an intensity curve that matches the experimentally observed curve Quantized: intensity goes towards zero at high frequencies Classical: intensity goes to infinity at high frequencies Bohr Model of Hydrogen -electrons move between orbits by accepting or emitting energy -can use equation r=n a/Z (z=proton#) to find radius of each orbit (n= energy level) Electron normally in n=1 (ground state), to remove electron it must move to n=infinity  called ionization -Electronic Transitions (absorbs or emits photons) -Electronic excitation: electron must absorb energy to jump UP levels (absorbs) -Electronic relaxation: electron must release energy to drop DOWN levels (emits) When calculating energy needed for electrons to jump levels: -Lyman series: bottom: n=1 -Balmer series: bottom: n=2 -Pachen series: bottom: n=3 de Broglie proposed that matter particles can be matter wave -he shot electrons through a double slit screen as evidence -wavelengths of matter are too short to detect hence why we don’t notice We can’t know the position and velocity of a particle exactly -error always exists (on formula sheet) Standing Waves Amplitude fixed at walls, troughs and crests occur at fixed positions and they include nodes which have an amplitude of zero -will not find a particle at a node Wave function tells you probability of finding electron at certain position Quantum Numbers Each unique set of n, l, ml and ms gives a different wave function (orbital) n = principle quantum number (row number) -any positive integer, determines energy of orbital, distance of electron from nucleus L = angular momentum quantum number -positive integers from 0 to n-1, determines shape of orbital L=0=s, L=1=p, L=3=d, L=3=f ml = magnetic quantum number -integer values from –l to +l, determines orientation of orbital ms= electron spin (axial rotation) -can either be -1/2 (beta) or +1/2 (alpha) Principle shell: contains n subshells and n^2 orbitals Subshell: contains 2l+1 orbitals -degenerate orbitals: have the same energies but differ in shapes Orbital Shapes  on paper Pauli exclusion principle: 2 electrons in a system can’t have same set of quantum # -when doing spdf notation; superscript number is number of electrons in the orbital -Cr(4s1-3d5) and Cu(4s1-3d10) are exceptions to filling rules Nodes Has n-1 total nodes L angular nodes n-1-L radial nodes Periodic Table Number of protons = Number of electrons =Atomic Number -elements in the same column have similar chemical properties (valence electrons) Atomic Radius: increases as you go down and right -Decreases with higher effective nuclear charge -Higher Z = smaller radius (if n is the same) -nucleus and electrons attract, electrons – electrons repel Ionic radii: more protons = smaller radii Two ions have same configuration: ion with higher nuclear charge is smaller -cations are smaller than parent atoms (2+ charge larger than 1+) -anions are larger than parent atoms Ionization Energy: decreases as you go down and right Energy required to remove electron from atom in gaseous form -first ionization: energy required to remove first electron in neutral atom (weakest) -second ionization: removal os 3nd electron from atom with 1+ charge, requires more energy than first ionization (group 13 and 16 violate this) -larger Zeff = large ionization energy -larger n = smaller ionization energy ElectronAffinity: decreases as you go down and right Amount of energy emitted when electron is added to atom in gaseous state -high electron affinity = more negative value -increases with Zeff and decreases with radius -after gaining first electron, energy for each electron gain becomes larger (group 2, 15 and 18 violate this) Effective Nuclear Charge: increases from left to right -stays somewhat constant as you go down a group Difference between true nuclear charge (Z) and effective nuclear charge is screened out by the core electrons (S) Zeff is the amount of positive charge on the nucleus perceived by electron Magnetic Properties Diamagnetic: all electrons are paired Paramagnetic: has unpaired electrons (odd number of valence electrons) **In all cases, when a full valence shell is attained the atom is very stable, any further changes require a lot of energy** Unit 2- Chemical Bonding Lewis Symbols Atomic symbol  Valence electrons as dots around symbol Exndnded Valences -2 row elements never exceed the octet rule Elements in row 3 or lower -Expanded valence table on sheet Transition metals can form up to sextuple bonds -atoms with n less than 3 can move left and right on table but must stay within the standard valences Formal Charges Charges assigned to individual atoms in a molecule -sum of formal charges must equal molecular charge formal charge= # of normal valence e – (# non bonded e + bonded e/2) Formal charges should only be -1, 0 or +1 -if atom moves left on table then +1 charge, right = -1 charge Resonance Structures -more than one lewis structure is correct -atoms do not move, bonds and electrons move -can only be resonance structures if all structures have the same formal charge Molecular Orbitals Tell us how electrons are distributed in molecules Sigma bond: orbitals overlap end-to-end and reinforce each other Sigma anti-bond: orbitals overlap end-to-end and cancel each other out Pi bond: orbitals overlap not end-to-end and reinforce each other (excluding s orbitals) Pi anti-bond: orbitals overlap not end-to-end and cancel each other (excluding s orbitals) -pi bonding orbitals are px and py orbitals -bonding orbitals: no nodes between atoms -anti bonding orbitals: at least one node between atoms Rules for molecular orbitals 1. number of molecular orbitals (MOs) = number of atomic orbitals (AOs) combined 2. two MOs formed from two AOs, 1 MO is bonding (lower energy) and 1 MO is anti- bonding (higher energy), electrons go into lowest energy MOs first 3. max number electrons in one MO is 2, MOs with same energy must follow Hund’s Atoms in the 2 Period 8 atomic orbitals are combined  1s and 3p + 1s and 3p which = 8 MOs -we get 4 sigma orbitals and 4 pi orbitals -we get 4 bonding orbitals and 4 anti-bonding orbitals If atoms are in group 13-15, the pi 2p orbital has lower energy than sigma 2p If atoms are in group 16-18, the sigma 2p orbital has lower energy than pi 2p 2s  *2s  2p  2p  *2p  *2p -remember both 2p orbitals can hold 4 electrons General Concepts -Bonding MOs formed when two AOs combined are lower energy than anti-bonding -Ordering of sigma2p and pi2p depends on atomic number -Anti-boding sigma2p is always least stable of the 8 MOs Triple bond: 1 sigma and 2 pi bonds -sigma bonds are stronger than pi bonds -sigma: side-to-side over lap, pi: on top overlap The bonding MOs are lower energy than theAOs The anti-bonding MOs are higher energy than theAOs Molecular Shapes VSEPR theory Molecule adopts geometry that maximizes distance between electron pairs -each lone pair and each bond (including multiple bonds) is one electron group -bond angles are determined by repulsion between electrons -lone pairs present decrease angle values between bonds Bond Order: Single =1, Double=2, Triple=3 -more electrons present increase the bond order, higher bond order = stronger bond bond order =(# of bonding electrons – number anti bonding electrons) /2 Bond Lengths: Distance between two nuclei of a bond -decreases with increasing bond order Polar bond = electrons are unshared -if electronegativity difference is from 0.5-1.7 then molecule is polar Unit 3- Molecular Shapes Dipole Moment -dipole moment exhibited by molecules which have a charge separation, units: 1 Debye -represented by a vector pointing towards more negative atom Steps to follow: 1. Draw the Lewis structure 2. Draw the vectors towards the more electronegative atom 3. Draw the VSEPR structure 4. Add bond dipoles and determine if they cancel out or not -Large dipole moments polar molecules -Small dipole moments  non-polar molecules When two atoms come together to form a bond… -bond length: distance between atoms nuclei at the potential energy minimum -bond length decreases with increasing bond order -sum of covalent radii of bonded atoms -bond energy: energy required to break the bond (always positive) -bond order: measure of number of electrons shared between atoms (double=2, triple = 3) Bond Energy Diatomic: 2H → H ΔH = -D(H-H) = -435.93 kJ/mol (g) 2(g) Multi-atomic: -average bonds in different scenarios -OH bond H2O  OH + H ΔH=D(H-OH) = 497.7 kJ/mol OH  O + H ΔH=D(O-H) = 428.0 kJ/mol Take the average and you getΔH= 463.4 kJ/mol ΔH  enthalpy  bond energy Chemical Bonding Hybridization: promoting an electron from s and an empty p orbital -s and p orbitals combine to make sp3 orbitals -total energy doesn’t change We can also combine an s orbital and just 2p orbitals to make sp2 orbitals -this leaves one normal p orbital -can
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