Laboratory II: Modeling Molecular Structure Part I

Purpose: Learn about the various structure and chemical bonding patterns of a variety of compounds. Specifically:

Lewis structures, hybrid orbital's, VSEPR, bond/molecular polarity,

resonance and chemical nomenclature.

Laboratory II

Commonality: Each model HELPS to represent and understand characteristics of molecules.

• Model 1-Ball and Stick: Sense of the way atoms are arrayed in molecules and of the molecule's three-dimensional structure– Why is 3D structure important?

• Molecular geometry: the three-dimensional arrangement of the atoms that constitute a molecule.

– Molecular Geo: helps determine several properties of a substance including its reactivity, polarity, phase of matter, color, magnetism, and biological activity.

Laboratory II

• Issue? Why not stop with Ball and Stick?– It is essential for a model to extend beyond a

simple appreciation of a static object but give us some sense of the object's potential.

• This model building exercise is designed to show you the value of molecular modeling structures have and to better inform you on molecular behavior and the limits of each model.

Model II: Lewis Dot Structure- Representation using Lewis symbols of covalent

bonding in a molecule• Diagrams that show the bonding

between atoms of a molecule, and the lone pairs of electrons associated with the molecule.

– Shared electron pairs are shown as lines

– Unshared electrons are show as lone or paired dots

• ONLY valence-shell electrons are shown

• Lewis Model Uses Rule of Orbital's, Octet Rule, and the Principle of Electroneutrality

The Rule of Orbital's: the total number of lone pairs and bond pairs (LP+BP) associated with an atom cannot exceed the number of Valence Shell Orbital's (VSO)

• VSO = n2, where n is the row of the Periodic Table in which that atom resides

Examples:

Hydrogen- n = 1: max VSE pairs (LP+BP) = VSO = 1

(B, C, N, O, F)- n = 2 : max VSE pairs (LP+BP) = VSO = 4 ("octet rule")

Model II- Lewis Dot Structure

Octet Rule: Atoms tend to bond in such a way that the atoms posses or share a total of 8 valence shell electrons

– Gives atoms the same electronic configuration as a noble gas.

– s2p6

• The quantum theory is used to explain the valence electron energy levels.

– Molecules tend to be most stable when the outermost electron shells of their constituent atoms contain eight electrons.

– Explains the stability and un-reactivity of noble gases (except helium)

Principle of Electroneutrality: Each atom in a covalent molecular assembly has a formal charge as close to zero as possible.

• Formal Charge (FC): The Charge of any atom in a molecule is the charge the atom would have if all the constituent atoms had the same electronegativity.

Rules and Theories help explain placement and rearrangement of electrons and formation of bonds.

Laboratory II: Steps for Lewis Structure

1.) Sum the Valence electrons. • If you have any cations (+) subtract them

from the total VE (group #)

• If you have anions (-) add them to the total VE.

2.) Write the central atom then the attached atoms.

• The central atom is generally the least electronegative atom (not including hydrogen).

• Electronegativity increases in the periodic table as you go from left to right and as you go up.

EXAMPLE: HCN

Step 1: H (1), C(4), N (5)

1+4+5-0 = 10 VSE 5 VSE pairs

Step 2a: C

Step 2b:

Laboratory II: Steps for Lewis Structure Cont.

3.) Complete the octets around the atoms bonded to the central atom.

• Make ‘Happy atoms’

EXAMPLE: HCN

Step 3: The two bonds use up four electrons so 6e- remaining

a.) Start w/minimal bonds

b.) Complete Octets on atoms around C-atom

Laboratory II: Steps for Lewis Structure

4.) Place any and all leftover atoms on the central atom, even if doing so violates the ‘normal’ octet rule.

• If there are not enough electrons to

give the central atom an octet, borrow e- from the adjacent molecules to form multiple bonds

Example HCN• The C-atom has 4e-, but according

to Step 4, needs 8e-

– What to do? Rearrange e- to make bonds.

Rearrange e-

Carbon Atom is still not Happy! Make more bonds!

Happy Molecules!

Laboratory II: Important Concepts of Lewis Structures

Most of the atoms in this exercise obey the “octet rule”

• Atoms with an expanded octet can accommodate more than eight electrons. – Atoms in the third and higher

periods of the periodic table– Transition metals– Representative groups (IIIA-

VIIA)

o Initially, assume that all heavy atoms obey the octet rule

o We will discover exceptions.

Exceptions and Violation of Octet Rule- When to Expand-

• Molecules and polyatomic ions containing an odd number of electrons

• Molecules and polyatomic ions in which an atom has fewer than an octet of valence electrons

• Molecules and polyatomic ions in which an atom has more than an octet of valence electrons

• The number of valence electrons is odd so a complete pairing of these electrons is impossible, in other words an octet around each atom is not possible.

Laboratory II Procedure

• Build ‘Ball and Stick’ of 16 molecules and deduce name, Lewis structures, geometry, polarity, resonance, and possible isomerism of each molecule.

• Work in pairs to build molecules together; DO NOT copy!– Work together to determine the answers

• The models you will use consist of wooden balls (atoms), pegs (single bonds) and springs (used for multiple bonds). – The balls represent the nuclei and inner core electrons. – The pegs and springs represent valence shell bonding or nonbonding

(lone pair) electron pairs.

Laboratory II: Concepts to Note

Resonance: Individual Lewis Structures in cases where 2 or more Lewis Structures are relatively equally adequate descriptions of the same Molecule.

• Structures must be stable- can not break rules

• The resonance structures in such instance are averaged to give a correct description of the real molecule. – Blend of all structures

Difference between Resonance and an Isomer

What does a Unicorn, Dragon, and Gryphon all have in common?

ResonanceThey are NOT REAL depictions

of the molecule.

• The real structure is a blend, compilation, an average, some intermediate of all the resonance structures

– Do not get caught up on static representation

• Just moving electrons to make or break bonds!

Resonance structures help explain molecule’s stability and/or reactivity

Example: Benzene Ring

Resonance

Be Reasonable• Follow the octet and

expanded octet rules

• Make molecules with Formal Charges close to zero as possible

• Satisfy VSERP

BAD Representation!!!!!

** Isolated Formal Charges** Octet not satisfied for Nitrogen atom

Structural Isomer

Isomer: Compounds whose molecules have the same overall composition but DIFFERENT STRUCTURES

• If it is possible to arrange the atoms of your molecule in another skeleton structure that is different from the one already made than it is an isomer

– Same molecular formula but different structural formula

Molecular formula: Chemical formula-describes what atoms are present in the molecule Example: Formaldehyde- CH2O

Structural formula: graphical representation of the molecular structure, showing how the atoms are arranged

Formaldehyde

Work an Example TogetherLab Example: Formaldehyde CH2O

Step 1: Draw Lewis Dot Structure

Step 2: Examine the Molecule's Molecular Geometry and describe the general overall shape of the molecule.

– Atoms linear? Triangle Shape?

– Use your textbook and black binder to help with geometry

Formaldehyde is flat (all of its atoms lie in one plane) with its peripheral atoms forming the vertices of a triangle.

Structure is called trigonal planar

Polar bonds are formed when two atoms with different electronegativity bond.

The more electronegative atom will get more than its fair share of electrons, and be the negative end of the bond.

The less electronegative atom will have some electron density pulled away, so it will be the positive end of the dipole.

Step 3: Predict the Molecule's Polarity

Step 3: Predict the Molecule's Polarity

Total dipole moment of the molecule is the vector sum of the bond dipoles. – If dipoles oppose one another

they will cancel out

From the shape and EN predict whether the molecule as a whole should be polar. – If you believe it to be polar,

indicate the direction of the molecule's dipole moment on the Lewis Dot Formula you drew for formaldehyde

Formaldehyde: The difference in electronegativity between C, at 2.5, and H, at 2.1, is not sufficient to give a significant bond dipole.

The difference between C, at 2.5, and O, at 3.5, is markedly polar.

Step 4. Examine the Molecule's Potential for Resonance

• There is only one place for the double bond in formaldehyde.

• There is no place for a triple bond to form

– Conclusion: No stable resonance structures.

Step 5. Examine the Molecule's Potential for Structural Isomerism

• If your molecule has structural isomers, build a model of at least one of them and compare it with your original model.

• It is possible to arrange the atoms of formaldehyde

– Build Molecule

– Deduce Formal Charge

– Stable?

Methanal: chemical formula: CH2 O

Formal Charges: Valence e- - (1/2 bonding e- + nonbonding e-)

Oxygen: 6- (1/2 *8 +0) = +2

Carbon: 4- (1/2 *4 +4)= -2

Due to separation of formal charges on the O and C atoms, there is no isomer

Molecules to Explorea. Methane (CH4):

– Construct a ball and stick model

• Note the symmetry

– Discern saturated or unsaturated: Saturated

– Note Geometry: tetrahedral molecule

– Carry out the additional structural analyses described in Part B above, like we did for formaldehyde

b. Dichloromethane (CH2Cl2) : – Convert to CH2Cl2 by replacing two

hydrogen atoms with chlorine atoms.

– Do the additional structural analyses described in Part B above.

– Draw the Lewis dot structures for all of the structural isomers of CH2Cl2 that you think may exist.

– Build structures of all of your isomers

• Compare them; are they super-imposable?

Will not adequately describe symmetry

Better representation Better representation

Molecules to Explore Cont.

c. Ethyl alcohol (ethanol): CH3CH2OH)

d. Dimethyl ether: (CH3OCH3)

• Construct models for ethyl alcohol and dimethyl ether.

• These two compounds are structural isomers, but they have very different properties.

• The differences are caused by the presence of the O-H bond.

The hydrogen bonding orientation is an important contribution to the forces between molecules in ethanol.

Ethanol Key Points: Very soluble in waterBp: 78.5°C

Dimethyl ether Key Points: Not water soluble Bp: -24°C

Assignments Due Date

• What is due: the worksheet with 16 molecules completely answered

• When is it Due: Hand into me by 4pm tomorrow

• Grade: based on correct answers, completeness, and turned into me on-time.

• Notebook Check: make sure to put any example molecules in your notebook– Formaldehyde

– Methane

– Dichloromethane

– Ethyl alcohol