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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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