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