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Chemistry 202

Imperial Valley College

IVC Supplemental Labs


Contents


M B (+packet) 2

Lewis Structures 2

M19 11

VSEPR Theory and Orbital Hybridization 11

IVC-14 13

Organic Chemistry-Physical Properties: Melting Point 13

IVC-15 14

Organic Chemistry-Chemical Properties 14

ADDITION TO: Experiment 15, General Chemistry in the Laboratory, Postma et al. 27

EXPERIMENT 15 27

ADDITION TO: Experiment 20, Polymers, Postma et al. 28

EXPERIMENT 20 28

ADDITION TO: Experiment 21, Heat of Vapor/Fusion, Postma et. al. 33

EXPERIMENT 21 33

ADDITION TO: Experiment 22, General Chemistry in the Laboratory, Postma et. al. 35

EXPERIMENT 22 35

Addition to M24: Determine the Equilibrium Constant for a Chemical Reaction 36

Purpose 36

Background 36

Materials 37

Equipment Setup 38

Procedure, Part 1 38

Procedure 38

Analyze 39

ADDITION TO: Experiment 25, General Chemistry in the Laboratory, Postma et. al. 43

EXPERIMENT 25 43

ADDITION TO: Experiment 26, General Chemistry in the Laboratory, Postma et. al. 44

EXPERIMENT 26 44

ADDITION TO: Experiment 36 (Exp 39 in 5th ed), General Chemistry in the Laboratory, Postma et. al. 46

EXPERIMENT 36 46

IVC-21 Kinetics of the Reaction between Iodide Ion and Peroxodisulfate Ion 47


The Iodine Clock Reaction 47

IVC-22: Voltaic Cells 54

IVC-23 Electroplating Galvanometric 59

IVC-24: Nuclear Radiation and Geiger Counters 67

Nuclear Chemistry 68

IVC-25 Determination of Δo in Cr(III) and Fe (III) Complexes 72

Cr(en)3 Cr(acac)3 & Fe(ox)3 synthesis 72

Spectrochemical series 72

GLX for UV Experiments 80

Equipment Setup 80

Procedure, Part 1 80

Laboratory portfolio 82






M B (+packet)




Lewis Structures



HANDOUT

G. N. Lewis, 1916 published a paper in JACS introducing the theory that a bond between two atoms resulted from the sharing of 2 electrons. (Lewis was from Berkeley and is also known for his theory of acid-base chemistry.) Lewis introduced this topic to his students to help them visualize how bonding produced molecular shapes.


Lewis Structures are based on the idea that:

  1. A bond consist of sharing 2 electrons

  2. Atoms with partially filled valance strive to fill their valance shell



A simple example of the first idea “A bond consist of sharing 2 electrons” is illustrated below between two p-orbitals atoms.





To describe the second idea “Atoms with partially filled valance strive to fill their valance shell” a new term Octet-Rule is introduced. This rule applies to nonmetals usually in the first two periods, and roughly applies to the remaining 4 periods as well.

  1. Atoms continue to form bonds until all their vacant orbitals are filled. Octet refers to the number 8, which is the maximum number of electrons it takes to fill ns2 np6 period.

  2. Once the ns2 np6 orbitals are filled the atoms has the same number of electrons as the Noble or Inert Gasses. The word Inert doesn’t necessarily imply unreactive, more then it implies containing the maximum number of electrons.



An example of fluorine: 19F 2s2 2p5


Fluorine is one electron from having a full octet. Adding a second electron fills the octet for Fluoride. No more electrons can be added to fluorine, fluorine will NOT continue to form bonds.





RULES FOR WRITING SIMPLE LEWIS STRUCTURES


l. Write the chemical formula for the molecule (or ion) and determine the total number of valence electrons in the molecule.


2. Draw the skeletal arrangement of the molecule showing single bonds connecting the atoms.


3. Assume that each bond in the skeleton requires two valence electrons (an electron-pair bond). After subtracting two electrons for each bond from the total number of valence electrons, assign the remaining electrons to give each atom an octet, or share of eight electrons.


4. If after each atom has been given a share of eight electrons, additional electrons remain, assign the extra electrons to the central atom of the molecule.


5. If there are not enough electrons to give each atom a share of eight electrons, then form multiple bonds between atoms by moving electron pairs to form double (or triple) bonds.


Examples of these rules with PF3, PO4-3 and NO3-


l. Write the chemical formula for the molecule (or ion) and determine the total number of valence electrons in the molecule.




Valance Electrons for P

Valance Electrons for F

Total Valance Electrons

PF3:

1 5

3  7

26


PO4-3: (1 5) + (4 6) + (3  1) = 32

NO3-: (1 5) + (3 6) + (1  1) = 24


2. Draw the skeletal arrangement of the molecule showing single bonds connecting the atoms.





3. Assume that each bond in the skeleton requires two valence electrons (an electron-pair bond). After subtracting two electrons for each bond from the total number of valence electrons, assign the remaining electrons to give each atom an octet, or share of eight electrons.





4. If after each atom has been given a share of eight electrons, additional electrons remain, assign the extra electrons to the central atom of the molecule.





5. If there are not enough electrons to give each atom a share of eight electrons, then form multiple bonds between atoms by moving electron pairs to form double (or triple) bonds.





Resonance Structures.





Possible Resonance Structures for NO3-





Formal Charge: Appling to each atom


Formal Charge = Group Number – Number of Bonds – Number of Unshared Electrons.

PF3




P: 5 – 3 – 2 = 0 P = 5 – 4 – 0 = +1 N = 5 – 4 – 0 = +1

F: 7 – 1 – 6 = 0 O = 6 – 1 – 6 = -1 O- = 6 – 1 – 6 = -1

O= = 6 – 2 – 4 = 0

Overall Charge = 0 -3 -1


How YOU draw a 3-D Shape



Electron Pair Shape Name (EPSN)

Molecular Shape Name (MSN): 0 lone pairs

Molecular Shape Name (MSN): 1 lone pairs

Molecular Shape Name (MSN): 2 lone pairs

Molecular Shape Name (MSN): 3 lone pairs

Linear (2)

Linear










Trigonal Planar (3)

Trigonal Planar










Trigonal Planar (3)




“V” bent







Tetrahedral (4)

Tetrahedral










Tetrahedral (4)




Trigonal pyramidal







Tetrahedral (4)







“V” bent




Trigonal Bipyramidal (5)

Trigonal Bipyramidal










Trigonal Bipyramidal (5)




Seesaw







Trigonal Bipyramidal (5)







“T”




Trigonal Bipyramidal (5)










Linear

Octahedral (6)

Octahedral










Octahedral (6)




Square pyramidal







Octahedral (6)







Square planar




Octahedral (6)










“T”

VSEPR


Valence shell electron-pair repulsion theory

To derive the VSEPR Class and VSEPR Name you must first find the Lewis Structure.








Bond Polarity: based on Electronegativity

Pauling ElectroNegativity scale


Linus Pauling devised a scale of electronegativity based on the relative ability of a bonded atom to attract the shared electrons.


Pauling gave Fluorine a EN value of 4.0 and Lithium 1.0. All other atoms fall between 0.8 and 4.0. The element with the smallest EN value is Cesium and Francium = 0.8.


H 2.1






















Li 1.0

Be 1.5

….

B 2.0

C 2.5

N 3.0

O 3.5

F 4.0

Na 0.9

Mg 1.2

….

Al 1.5

Si 1.8

P 2.1

S 2.5

Cl. 3.0

K 0.8

Ca 1.0

….













Br 2.8


For reference: Period 2 elements Fluorine, 4.0, has the GREATEST ability to Attract Electrons. Li has the LEAST ability to Attract Electrons.


i.e. C-F, C-Cl, C-Br, C-I….rank from most to least EN.


Two types of bonds

  1. Ionic---metal & nonmetal

  2. Covalent---nonmetal & nonmetal

Bonds are further classified based on electronegativity.

  1. Ionic (transfer of electrons)

  2. Polar covalent (unequal sharing of electrons)

  3. Nonpolar covalent (equal sharing of electrons)

Examples of Polar Covalent using Partial Positive + and Partial Negative -





Which is a Polar/Nonpolar: SO3 -vs- SO2





Symmetrical Unsymmetrical


Polarity is determined by:

  1. Electronegativity charge calculation

  2. Symmetry of molecule


Example of polarity: NF3 is polar because it’s unsymmetrical with Lone Pairs. CH2Cl2 is unsymmetrical because if has different atoms, so it is polar. SF4 is polar because it is unsymmetrical because of the Lone Pairs.





For Your Lab Manual: Write these 7 columns…ACROSS TWO PAGES of your lab manual. Include all items listed under each column


M B (+packet) Example Lewis Structures

 Formula

Number of Valance Electrons

Lewis Structure with Formal Charge

VSEPR electron pair shape name for each Central Atom, Plus Oxidation number

For each Central Atom, and Overall molecular shape find the VESPR Molecule shape name, draw the 3-D Shape and Bond Angles

3D structure with Polarity Arrows then answer: is it polar = Yes or No

Resonance Structure

CF4

C = 4

F = 4@7=28

Total = 32



AX4

Tetrahedral

Tetrahedral


No: symmetrical


No

NH3

 

 

 

 

 

 

 

I will grade EACH Page, do not go to the second page until I've graded the previous page.

 


Extra Lewis Structures…put these in your lab book.

C2H6, C2H4, C2H2, CO3-2, and CH3NH2


I want to see 24 structures.


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