Download Lewis Structure Lab and more Exams Geometry in PDF only on Docsity!
Experiment #2. Lewis Structure
A Lewis structure shows how the valence electrons are arranged and indicates the bonding between atoms in a molecule. We represent the elements by their symbols. The shared electron pair is shown as a line/bond between the two atoms. All the other valence electrons are shown as dots or lines around the symbol of the element.
For Example: The Lewis structure for Cl 2 is Let us now see how to draw the Lewis structure for CO 2
Steps for drawing Lewis Structures Example with CO 2
- Sum the valence electrons of all the atoms. For anions, add one electron for each negative charge. For cations, subtract one electron for each positive charge.
The total number of valence electrons is 16 (4 from carbon and 6 from each oxygen).
- Choose the least electronegative element (other than H) as the central atom. For a molecule of type XYn, X is always the central atom.
Choose carbon as the central atom since it is less electronegative and also follows the XYn rule.
- Use single bonds (lines) to connect the central atom to the surrounding atoms.
Attach the two oxygens to carbon by single bonds. O – C – O
- Subtract 2 electrons for each bond from the original total number of electrons.
We subtract 4 electrons (2 for each bond) from 16 leaving 12 electrons to distribute to the remaining atoms.
- Complete the octet for all outer atoms (other than H). Each bond to an atom will count as 2 electrons for that atom's octet.
Complete the octet for each of the oxygen atoms which uses the last 12 electrons. (We need to add 6 electrons to each oxygen.)
- Complete the octet for the central atom. There are no more electrons, so we cannot add
more to the central atom to complete its octet.
- If you run out of electrons before the octet for the central
atom can be completed, start forming multiple bonds until the central atom has a complete octet. When you form a multiple bond, remember to remove an electron pair from the outer atom.
Since carbon does not have an octet, form multiple bonds by taking the lone pairs.
becomes
The central atom is still electron deficient, so share another pair.
becomes
- Make sure the correct number of electrons has been used
and that every atom’s octet is complete. The exceptions are atoms of group 2 and 3 which can have incomplete octets and atoms in period 3 or greater which can have expanded octets.
4 bonds + 2 electron pairs equal 16 electrons. All atoms have an octet.
Molecular Geometry – the Valence Shell Electron Pair Repulsion (VSEPR) Theory The electron groups around the central atom repel each other and therefore prefer to be as far apart from each other as possible. This is the main idea of the VSPER theory. We can apply the VSEPR theory to predict the molecular shape/geometry of a molecule.
- Draw the Lewis structure for the molecule in question.
- Count the total number of electron groups on the central atom. Add the number of atoms bonded to the central atom and the number of lone pairs on the central atom – this is the total number of electron groups. Note that multiple bonds to one outer atom still count as one electron group.
- The arrangement of the electron groups is determined by minimizing the repulsions between them.
- Remember that lone pairs require more space than bonding pairs. Therefore, choose an arrangement that gives lone pairs as much room as possible.
The attached table shows the relationship between the number of electron pairs and the molecular geometry.
Polarity of a molecule A covalent bond is polar if there is a difference in electronegativity between the bonded atoms. A molecule like HCl has a polar covalent bond since there is a difference in electronegativity between the two atoms (the difference is greater than 0.4 and less than 1.8). Thus HCl possesses a permanent dipole moment because the molecule has a distinct negative end and a distinct positive end. However, just because there is a polar bond present in a molecule does not necessarily mean that the molecule is polar. If all the dipoles in the molecule cancel each other out, then the molecule will be non-polar. For example, CO 2 has two polar bonds, but they point in opposite directions and cancel each other out.
Formal Charges
The formal charge of any atom in a molecule is the representation of electron distribution on the atom. Remember that the formal charge does not represent the real charge on the atom. It is a fictitious charge assigned to each atom that helps in finding the best Lewis structure for a molecule.
The best Lewis structure will Have the lowest possible formal charge on each atom Put the negative formal charge on the most electronegative atom (and a positive formal charge on the least electronegative atom)
Calculating formal charges
- Sum all the electrons in the lone pairs belonging to that atom
- Add to this half of the bonding electrons
- Subtract this total from the number of valence electrons for that atom to get its formal charge.
In the above example for CO 2 , the Lewis structure on the Left is the better one as it places a formal charge of zero on each atom. Recall that the formal charges sum to zero for a molecule and to the charge for an ion.
VSEPR Geometries
Number of electron groups on the central atom
Hybridization
0 lone pair on central atom
1 lone pair on central atom
2 lone pairs on central atom
3 lone pairs on central atom
4 lone pairs on central atom
2 sp
Linear
3 sp^2
Trigonal planar (^) Bent
4 sp^3
Tetrahedral Trigonal pyramidal Bent
5 sp^3 d
Trigonal bipyramidal See-saw / sawhorse T-shape Linear
6 sp
3 d 2
Octahedral (^) Square pyramidal Square planar^ T-shape^ Linear
Name_________________________
CHM112 Lab – Lewis Structures – Grading Rubric
Subject to additional penalties at the discretion of the instructor.
Criteria Points possible Points earned
Question 1 (1.33 points each question) 8
Question 2 (1 point each question) 4
Question 3 (1 point each question) 2
Question 4 (1 point each question) 4
Question 5 (1 points each question) 2
Total 20
CHM 112: Lewis Structures Lab Name ________________________
- Fill in the table below.
Lewis Structure (redraw to reflect geometry, show net dipole direction if polar)
Polar? Y/N Geometry
CH 2 I 2
CBr 4
CBr 2 O
NCl 3
CH 3 CH 2 OH at C?
at O?
C 5 H 10 O
(connect C’s & O in a 6- membered ring. H’s are attached to C’s )
at C?
at O?
- Draw the Lewis Structures of each of the following acids: (These structures, like most acids, have each H attached to
one of the oxygens.
(A) H 2 SO 4 (B) HIO 2
(C) H 3 PO 3 (D) CF 3 CO 2 H (Both O atoms are attached to second C)
- Draw the Lewis Structure for these anions. Each anion is a conjugate base of an acid.
(E) IO 3 ─^ (F) SO 3 2─
- Weak bases frequently contain N: (nitrogen with a lone pair of electrons). Draw the Lewis Structure for each of the following weak bases.
(A) NH 3 (B) CH 3 CONH 2 (Connectivity is )
(C) CH 3 NH 2 (D) C 5 H 5 N (six membered ring made of C and N, Hs attached to
Cs)
- Draw the Lewis structures of the following cations. Each cation is the conjugate acid of a weak base:
(A) NH 4 +^ (B) CH 3 NH 2 CH 3 +
