1. Atomic and Molecular Structure
Michael Abosch
Brian Pflaum
2nd Period

2. Unit Outline Atomic and Electronic Structure, and Quantum Mechanics
Periodic Trends
Molecular Structure
Bonding Theory

3. The Wave Nature of Light
Electromagnetic Radiation- all visible light, radio waves, infrared, X-rays etc.
Electromagnetic Spectrum- shows radiation arranged in order of increasing wavelength
Visible light is only a small portion of spectrum.

4. The Wave Nature of Light
fλ= c
(frequency)(wavelength)= Speed of light (2.9979x108 m/s)
Frequency measured in s-1 (often Hz)
Wavelength measured in meters (often nm,μm)

5. The Quantization of Energy
Quantum=The smallest quantity of energy that can be emitted or absorbed as electromagnetic radiation.
Energy, E, of a single quantum equals a constant times the frequency of radiation.
E=hf
h=planck’s constant=6.626X10-34Joule-seconds.

6. Photoelectric Effect
When photons of sufficiently high energy (greater than the individual metal’s threshold energy) strike a metal surface, electrons are emitted from the metals
Energy of Photon, E=hf (planck’s constant)(frequency)
Kinetic Energy of ejected electrons: KEe=Ephoton-Ethreshold of metal

7. Wave Behavior of Matter
Dual nature of radiant energy: both particle and wave-like properties
DeBroglie wavelength: wavelength=(Planck’s constant)/(momentum)=(h)/(mv)
Mass in Kg, Velocity in m/s

8. Orbitals
An allowed energy state of an electron in the quantum mechanical model of the atom; describes the spatial distribution of the electron. The orbital is defined by the values of quantum numers n, l, and ml

9. The Principal Quantum Number
The principal quantum number, n, can have positive integral values of 1,2,3 etc…As n increases, the orbital becomes larger, and the electron spends more time farther from nucleus

10. The Azimuthal Quantum Number
The azimuthal quantum number, l, can have integral values from 0 to n-1 for each value of n. This quantum number defines the shape of an orbital. The value of l is generally designated by the letters s,p,d, and f.
Value of l 1 2 3
Letter Used s p d f

11. The Magnetic Quantum Number
The magnetic quantum number, ml, can have integral values between -l and l. Describes orientation of orbital in space.

12. Relationship amongst Quantum Numbers

13. Spin Magnetic Quantum Number and Pauli Exclusion Principle
The Spin Magnetic Quantum Number, ms, has two possible values: +1/2, -1/2.
No two electrons in an atom can have the same set of four quantum numbers n, l, ml, and ms
Thus, an orbital can hold a maximum of two electrons, and they must have opposite spins.

14. Electron Configurations
Electron Configuration=A particular arrangement of electrons in the orbitals of an atom.
The orbitals are filled in order of increasing energy, with no more than two electrons per orbital.

15. Orbital Diagrams
Each orbital is denoted by a box, and each electron by a half arrow (which represents spin-up or spin-down)
Electrons having opposite spins are said to be paired when they are in the same orbital
An unpaired electron is one not accompanied by a partner of opposite spin.

16. Hund’s Rule
Hund’s Rule=For orbitals of the same energy level, the lowest energy is attained when the number of electrons with the same spin is maximized.
Note how in the diagram below, all three p orbitals are filled singularly before an electron is paired

17. The Periodic Table and Electron Filling Order

18. Condensed Electron Configurations
The Electron configuration of the most recent nobel gas is
represented by its chemical symbol in brackets. From there,
Just proceed in the normal filling order until you reach the element.
In Potassium, the previous noble gas is argon, and its remaining
Electron occupies just one of the s orbitals, hence why it is denoted
As 4s1

19. Ions
Start by writing the electron configuration for the normal element
Then remove (or add) electrons as necessary, always taking (or adding) from the highest principle quantum number first (ignoring the filling order).
Fe=[Ar]4s23d6
Fe(II)=[Ar]3d6

20. Anomalous Electron Configurations
Electron configurations of certain elements appear to violate the “rules”
Frequently occurs when there are enough electrons to lead to precisely half-filled sets of degenerate (same energy-level) orbitals, or to completely fill an orbital. This conserves Energy
No universal pattern or predictability
Ex: Chromium is [Ar]4s13d5 instead of [Ar]4s23d4

21. Practice What’s the electron configuration for Lead?
Answer: [Xe]6s24f145d106p2
Assign Quantum numbers to it’s last filled electron.
Answer: n=6, l=1, ml=0, ms=+1/2

22. Periodic Trends Atomic Size Ionic Size Ionization Energies
Electronegativity
campus.ru.ac.za/full_images/ img jpg

23. Atomic Size
Within each group, size increases from top to bottom, results primarily from the increase in principle quantum number of electrons
In each period, atomic radius tends to decrease from left to right. Increase in the effective nuclear charge as we move across a row steadily draws valence electrons closer to nucleus
Exceptions: The addition of a paired electron produces increased repulsion that sometimes leads to an increase in size (Like from a p3 to a p4 element.)

24. Atomic Size

25. Ionic vs. Atomic Size
Cations: Compared to its neutral atom, cations are smaller because electrons have vacated the biggest orbital
Anions: Compared to its neutral atom, anions are larger because adding electrons increases repulsions, which leads to more space.

26. Ionic vs. Atomic Size

27. Isoelectronic Series
Isoelectronic Series=A group all containing the same number of electrons. As the atomic number increases, the radius decreases.
Ex: Cl-, Ar, K+
Size: Cl->Ar>K+

28. Ionization Energy
Ionization Energy=The minum energy required to remove an electron from the ground state of the isolated gaseuous atom or ion
The Greater the ionization energy, the more difficult it is to remove an electron.

29. Variations in Successive Ionization Energies
I1>I2>I3 etc…
It’s more difficult to pull away an electron from an increasingly more-positive ion
There is a sharp increase in ionization energy to remove a core electron, as they are closer to the nucleus.

30. Periodic Trends in First Ionization Energy
Within each period, ionization energy generally increases with increasing atomic number.(Smaller atomic radius)
Within each group, Ionization generally decreases from top to bottom (Larger atomic radius).
Irregularities: Added “p” orbital sometimes leads to decrease in ionization energy because the “p” orbitals have more space than the “s” orbitals. Adding a paired electron can also lead to a decrease in ionization energy, as there is increased electron-electron repulsion.

31. Periodic Trends in First Ionization Energy

32. Electronegativity
Electronegativity=An order of an atom’s overall ability to attract electrons. It combines atomic size and ionization energy into a single summary number.

33. Covalent Bonding Created when two atoms share electrons
Strive to fulfill the Octet rule- “atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons”
Many covalent bonds are exceptions to the octet rule

34. Lewis Symbols
Consists of the Atom’s chemical symbol, plus one dot for every valence electron it has
Anions have extra dots, cations fewer dots
Examples:
H• : Ar : :F: • C •

35. Drawing Lewis Structures
Write the Chemical symbols for every atom in the molecule
The atom that makes the most bonds is generally the central atom
Determine the total amount of Valence Electrons in the molecule
Place single bonds between all atoms in the molecule that bond
With remaining electrons, fill up octets on all the atoms
If extra electrons exist, place them on the central atom
If too few electrons exist, create double, or triple bonds, keeping the octet rule in mind.

36. Drawing Lewis Structures
Write all Chemical symbols
Example- CO2
Carbon makes more bonds (4) than oxygen (2)
O+O+C = 6+6+4= 16
Place single bonds
O C O
- -
= =
Fill all Octets
Not Enough! Must make double bonds
This Creates 16 electrons, while satisfying the octet rule

37. Formal Charge
Formal charge= the charge the atom would have if each bonding electron pair were shared evenly between its two atoms
To determine formal charge draw Lewis structure, and
Count all unshared electrons per atom
Add half of the single, double, or triple bonds electrons to the total (either 1,2, or 3 electrons)
Subtract this number from that atom’s usual amount of valence electrons

38. [:C≡N:]- Formal Charge Example- CN- -1 Count all unshared electrons
Add half of bond total
[:C≡N:]-
Subtract from Atom’s usual amount of valence electrons 2 2
+(6/2)
+(6/2) = 5 = 5 -1

39. Electron Domains
Any Bond (only single bonds) plus electron pairs (or last unpaired electron) counts as an electron domain.
Electron Domains are important in understanding molecular shape
Shapes are categorized by the amount of total electron domains, then described further by the amount of bonding domains
If an atom has 5 electron domains, but only 3 are bonding domains, the other 2 are considered non bonding domains, and are lone pairs.

40. Molecular Shapes 5 Basic Shapes
Shape based on number of electron domains in the molecule
Linear
Trigonal Bipyramidal
Trigonal Planar
Octahedral
Tetrahedral
All Pictures: chemlab.truman.edu/.../MM1Files/Linear3.gif

41. Linear One or Two electron Domains 1 or 2 bonding domains
Bond angles = 180˚
Example- CO2

42. Trigonal Planar Three Electron Domains Bond angle = 120˚
3 bonding domains- trigonal planar
Ex. BF3
2 bonding domains- bent molecule
Ex. bent- NO2
Trigonal Planar
Bent

43. Tetrahedral Four Electron Domains Bond Angle109.5˚
4 bonding domains- Tetrahedral
ex. CH4
3 bonding domains- trigonal pyramidal
ex. NH3
2 bonding domains- bent
Ex. H2O
Trigonal pyramidal
Bent

44. Trigonal Bipyramidal Five Electron Domains
Bond Angles- Equatorial 120˚ Polar 180˚
5 bonding domains- trigonal bipyramidal- ex. PCl5
4 bonding domains- Seesaw-ex. SF4
3 bonding domains- T-shaped- ex. ClF3
2 bonding domains- Linear- ex. XeF2
Linear
See-Saw
Trigonal Bipyramidal
T-Shaped

45. Octahedral 6 Electron Domains Bond Angles- Equatorial- 90˚, Polar 180˚
6 bonding domains- Octahedral
Ex. SF6
5 bonding domains- Square Pyramidal
Ex. BrF5
4 bonding domains- Square Planar
Ex. XeF4
Octahedral
Square Pyramidal
Square Planar

46. Dodecahedral
Just Kidding

47. Bond Order & Length Double bond= bond order of 2
Triple bond = bond order of 3
As Bond order increases, bond length decreases
As Bond order increases, greater repulsive forces exist between adjacent electron domains, creating bigger angle
As Bond order increases, more energy is needed to break the bond

48. Bond Polarity Happens when electrons are shared unevenly between atoms
Therefore does not happen between like atoms (i.e. H-H)
Generally, electronegativity differences of .4 or higher are considered polar
When electronegativity difference is great enough, the bond is considered ionic, not polar covalent
Ex H-H C≡N Na-Cl
= = 2.1
0< > >>.4
Nonpolar Polar Ionic

49. Molecular Polarity When is a molecule Polar? No Yes Yes
Polar Bonds Present?
Polar Bonds arranged symmetrically?
Nonpolar Molecule
Polar Molecule

50. Molecular Polarity Symmetrical Molecules Asymmetrical Molecules Linear
Trig. Planar
Bent
Square Pyramidal
Tetrahedral
Trig. Bipyramidal
Pyramidal
Seesaw
T-Shaped
Square Planar
Octahedral

51. Valence Bond Theory (hybrid orbitals)
Bonds occur when electron shells overlap
Since electrons are simultaneously attracted to both nuclei, bonds occur
Valence bond theory alone does not explain polyatomic molecules. For this, Hybrid orbitals are needed

52. sp orbitals Consider the linear non-polar BeF2 molecule 1s 2s 2p
As it is, the molecule would not bond, since it has a full 2s shell
After promotion…
The molecule can now bond, however it would make non identical polar bonds
1s 2s p
By Hybridizing the 2s and 2p shells…
Now the Be molecule can make 2 identical bonds
1s sp p
The remaining 2p orbitals end up unhybridized
Additionally, the bigger lobes produced by the sp orbital allow for more overlap, which means stronger bonds
All 2p orbitals can be hybridized. When this occurs, the same amount of orbitals must be created. Ex.
1s 2s p promote 1s 2s p hybridize 1s sp p

53. More Hybrid Orbitals sp makes 2 180˚ orbitals sp2 makes 3 120˚orbitals
sp3 makes tetrahedron-arranged orbitals

54. Sigma(σ) and Pi(π) Bonds
σ bonds- bonds occurring on the internuclear axis
π bonds- bonds occurring between two p orbitals oriented perpendicularly to the internuclear axis.
π bonds produce a sideways overlap, which is not as substantial, and therefore, not as strong, as a σ bond
Single bonds are σ bonds, double bonds consist of one σ bond and one π bond, triple bonds have one σ bond and 2 π bonds

55. Molecular Orbital Theory
Better explains excited states of molecules
Each Molecular Orbital (MO) holds up to two electrons, of opposing spins
Associated with the entire molecule, not just a single atom

56. Molecular Orbital Theory
Easiest way to analyze is through an energy level diagram
Bottom half of each shell is the bonding molecular orbital, and is lowest energy.
Top half is the antibonding molecular orbital, and is higher in energy
As the name suggests, antibonding orbitals cancel out the bonding orbitals
Because energy increases as the position on the chart increases, slots are filled in from the bottom up

57. Molecular Orbital Theory
Additionally, based on the positioning in the diagram, the bonds can be analyzed as either σ or π bonds
The diagrams can be used to determine whether or not an atom would form naturally

58. Bond Order
Bond order = ½(# of bonding electrons- # of antibonding electrons)
Bond order of H2 = ½(2-0) = 1, Therefore H2 exists
Bond order of He2 = ½(2-2) = 0. Therefore, Helium is not diatomic in nature