by jake grodsky and sarine hagopian. image from:
TRANSCRIPT
UNIT VI: Atomic
and Molecular StructureBy Jake Grodsky and Sarine Hagopian
Atomic and Electronic Structure/ Quantum
Mechanics
Electromagnetic Radiation and Spectrum
Image From: http://www.eoearth.org/files/115601_115700/115629/350px-Spectrum.jpg
Atomic Spectra 1. An electron in the atom absorbs energy from heat,
electricity, radiation, etc.2. That electron moves to an orbital at a higher
energy level3. Later, the excited electron returns to a lower
energy level4. Excess energy lost by electron is released as light
or other electromagnetic radiation Since each element has its orbitals at slightly
different energies, each spectrum has a unique finger print.
Atomic Spectra
http://www.youtube.com/watch?v=QI50GBUJ48s
Energy Quantization
•According to Niels Bohr’s theory: electrons can only exist in certain possible energy levels.•Energy of an electron is proportional to its distance from the nucleus
Image From: http://hyperphysics.phy-astr.gsu.edu/hbase/imgmod/bohr1.gif
Bohr Model
Image From: http://reich-chemistry.wikispaces.com/file/view/rutherford_bohr_model.gif/103784023/rutherford_bohr_model.gif
Photoelectric Effect•When light is shone on metal, electrons are emitted from the metal
•The effect can be used to switch a light signal into an electric current
Bright light
Dim light
0 fthreshold
curr
ent KEejected e-=Ephoton-Ethreshold
Wave-Particle Duality
•For a long time it was believed that light was solely a wave •Both light and electrons have a dual nature •They exhibit characteristics of both waves and particles •The photoelectric effect proves that light has a particle nature as well•The wave properties of electrons are shown through the DeBroglie Hypothesis
DeBroglie Wavelength
λ = wavelength of a particle
Constant: me = 9.11 × 10-31 kg
velocity
Constant:h = 6.626 × 10-34 J•s
Wave-Mechanical Model
•The wave mechanical model is the most recent model of the atom
•Improvements were made on Bohr’s model, specifically dealing with electrons
•Electrons are treated as waves instead of particles- electron has more in common with
light, tv, radio waves, microwaves, and x-rays than it does with protons and neutrons
•Orbitals are the regions in atoms which are most likely to have electrons in them
•The model is more statistical than visual
•This model includes energy levels which are numbered 1-7(closest to farthest) which
indicates how far a given electron is from the nucleus
•The energy level can be viewed in the same way as Bohr’s model viewed the shell
Electron Configurations of Atoms and Ions
•Lower energy levels are always filled first•Ions have less electrons than the neutral parent atom• this means that
electron configurations of ions look like those of other neutral elements (even from a different atom)
Image From: http://www.mikeblaber.org/oldwine/chm1045/notes/Struct/EPeriod/IMG00011.GIF
•Configuration can be abbreviated • The last Noble Gas element symbol is put
in brackets and remainder of the electron configuration is written out
Example: Electron Configuration of Zinc: 1s22s22p3s23p64s2 3d10
Abbreviated Electron Configuration of Zinc: [Ar] 4s2 3d10
Image From: http://www.mpcfaculty.net/mark_bishop/abbreviated_electron_configuration_help.htm
Quantum Numbers
Image From: http://library.thinkquest.org/19662/images/eng/pages/improved-bohr-2.jpg
•The final quantum number is the ms number. This is ± ½, depending on the spin of the electron
Key:
n= principal quantum numberl= angular momentum number [0- (n-1]ml= magnetic quantum number [-l – l]
•Electron configurations can be expressed as orbital diagrams as pictured below by visualizing each individual electron and its corresponding spin, as well as the orbital and energy level that it is a part of.
Orbital Diagrams
Oxygen
Electron Configuration: 1s22s22p4
Total number of e-: 8
1s 2s 2px 2py 2pz
What would the quantum number be for this e-?
2, 1, -1, -½
•To form an orbital diagram:1. Determine the electron configuration of the atom and the total
amount of electrons.
2. Following Hund’s Rule, begin to fill orbitals from lowest energy level to highest, remembering the Pauli Exclusion Principal and having an upward and downward arrow in each orbital, representing the positive and negative electron spin (responsible for the +½ and -½ values of ms.
Periodic Trends
Atomic SizeAtomic Radius: Size of an atom which is influenced by the volume of the e- orbitals (clouds)
Decreases
Incr
ease
s
Why does atomic radius increase as you go
down a group?
• more energy levels so the new levels are
“blocked” and therefore not as tightly pulled to
the center
Why does atomic radius decrease as you go across a period?
• Only one p+ and one e- are added
•Increasing nuclear charge pulls outermost e-s closer and closer to the nucleus
reduces atom size
•All additional e- go into same principle energy level so shielding is not an issue
nucleus just gets stronger and squeezes everything closer
Ionic Size
•Cations are smaller than their neutral parent atoms
• Cations have more protons than electrons (hence their positive charge)
• Protons more tightly pull the electrons towards the nucleus therefore reducing the
size of the atom
•Anions are larger than their neutral parent atoms
• Anions have more electrons than neutrons (hence their negative charge)
• Electrons are not as attracted to the nucleus therefore increasing atomic radius
Ionization Energy and Electronegativity
Ionization Energy: amount of energy needed to remove an e- from an atom or ion
Electronegativity: a measure of the ability of an atom in a chemical compound to attract/gain e-s
Increase
Dec
reas
e
•As you go down a group, more orbitals
are added valence e- are farther from
nucleus so pull of p+s on the e-s is reduced
•As you go across a period, more protons
are added to the nucleus valence
electrons are held more tightly
Molecular Structure
Covalent Bonding•Formed when e- pairs are shared amongst atoms
•Generally a metal and a nonmetal pair
Key Terms
•Lewis structure: representation of a covalently bonded molecule and its valence electrons
•Octet Rule: In a covalent molecule, each atom has eight electrons around it
•Lone Pair: pair of e-s not involved in bonding
•Bond pair: pair of e-s shared between two atoms
•Double bond: two pairs of e-s shared between atoms
•Triple bond: three pairs of e-s shared between atoms
•As the number of shared e- pairs goes up, bond length goes down }
Lewis Structures
•Symmetrical arrangements are more likely than asymmetrical ones
•The less electronegative atom tends to be in the middle
•Subtract valence e-s from total electrons needed to complete octet/duet and divide by two this is the number of bonds that will need to be made
•If e- needed > e- remaining, add bonds
•If e- remaining > e- needed, add lone pairs to central atom
Image from: https://vinstan.wikispaces.com/file/view/lewis_structure.gif/46694367/lewis_structure.gif
Polarity
Are any bonds polar?
Are polar bonds arranged symmetrically?
Polar Molecule
Non-Polar Molecule
Yes
YesNo
No
•Electronegativity difference of:0.5 or less bond is nonpolargreater than .5 bond is polargreater than 1.7bond is ionic
Hybridization, Shape, and Angle Number of Electron Domains
Electron Geometry
Bonding Pairs
Non-Bonding
Pairs
Molecular Geometry
HybridizationBond Angle
2 Linear 2 0 Linear sp 180º
3Trigonal Planar
3 0Trigonal Planar sp2 120º
2 1 Bent
4 Tetrahedral
4 0 Tetrahedral
sp3 109.5º3 1Trigonal
Pyramidal
2 2 Bent
5Trigonal
Bipyramidal
5 0Trigonal
Bipyramidal
sp3d120º90º4 1 Seesaw
3 2 T-Shaped
2 3 Linear
6 Octahedral
6 0 Octahedral
sp3d290º90º
5 1Square
Pyramidal
4 2Square PlanarSquare Planar
Image from: http://www.chem.hbnu.edu.cn/jysweb/whjys/wangwd/jghxywkj/AX4E2.gif
SeesawImage from: http://www.chem.hbnu.edu.cn/jysweb/whjys/wangwd/jghxywkj/AX4E1.gif
Trigonal PyramidalImage from: http://www.chem.hbnu.edu.cn/jysweb/whjys/wangwd/jghxywkj/AX3E1.gif
Formal Charge• Helps decide which Lewis structure is most
reasonable• It is the charge the atom would have if all the
atoms in the molecule had the same electronegativity
• To calculate: 1. Count all nonbonding electrons per atom. 2. Count half of any bond. 3. Subtract valence electrons by the number assigned to each atom.
Bonding Theory
Valence Bond Theory (Hybrid Orbitals)
Image From: http://courses.chem.psu.edu/chem210/quantum/pictures/sigms.gif
Valence bond theory says
that electrons in a covalent
bond can be found in a
section that is the overlap
of the individual atomic
orbitals that are bonding
Molecular Orbital Theory
Sigma and Pi BondsSigma bonds = formed by the overlap of two s orbitals, an s and a p orbital, or two p orbitals
Pi bonds = formed by the overlap between two p orbitals oriented perpendicularly to the internuclear axis
Where do Pi bonds come from?
•Only period 2 elements form pi bonds because they are small in size and therefore form short bonds.
•Due to these short bonds when 2 of these atoms form a sigma bond, they are so close together, that an additional energy level(orbital) overlap and a pi bond is formed
Ex) O=O has 1 sigma bond and 1 pi bond N≡N has 1 sigma bond and 2 pi bonds
Bond Order
½ [ (#bonding e-) - (#anti-bonding e-)]
• E = hf
• KE = ½mv2
• c = fλ
• λ =
•c = 2.9979 × 108 m/s •h = 6.626 × 10-34 J•s
•me = 9.11 × 10-31 kg •NA = 6.022 × 1023
Equations and Constants