Notes:
HED stands for High Electron Density so that would be the same thing as Clouds.
This link will also help with Bond Angles:
http://en.wikipedia.org/wiki/Molecular_geometry

1 Lewis Dot Notation

· This question should read “in order of decreasing size.”
Arrange the ions Be2+, F−, Mg2+, S2−, Cl−
in order of increasing size.
· Radicals are species with an unpaired electron. Draw the Lewis dot for each.
Which of the species
NO, BrO, CH3+, BF4-
are radicals?

Smog

When NO2 and NO3 react you product N2O5

2 Lewis Dot Notation

· This question is tricky read the note:
Rank the following atoms and ions
Li+, Be2+, He, H−, B3+
in order of decreasing size.
These particles have the same number of electrons. The number of protons in these ion is the biggest determinate of the size of the particles, if the electron number is constant. Draw a diagram of each particle. Remember opposites attract.

Additional Notes

  • When an atom has a half filled or full sublevel there is sometims a discrepency in the general trends that you observe. Like it is easier to remove an electron from oxygen than nitrogen. Nitrogen has a haf filled "p" sublevel giving it more stability.

Resonance

If you have a question about resonance, the video 8.3 below is very useful. Watch it before you ask me.

Structure least energy

http://www.uwosh.edu/faculty_staff/gutow/Lewis_Tutorial/Lewis.html
If this link doesn't work, see the bottom of the screen. I have copied and pasted the material from this page.

Xenon Molecules:

Each xenon atom has 8 electrons plus it will add a single bond for each halogen and a double bond for each oxygen family
2015-01-06_1237.png

Video:


Watch this video to explain the O3 bonding:

(Watch these videos as needed. You may also want to look at the Bonding videos if you haven’t done that unit.) Answer these Questions

Types of Bonding Electronegativity
Bond Polarity Dipole Moments
e- configs
**8.1**
Forming Binary Ionic Bond Energy
Lewis Structures Octet exceptions__
**8.2**
Resonance VSEPR
**8.3**

Videos you may need:

8.4: Ions: Electron Configurations and Sizes
8.5: Formation of Binary Ionic Compounds
8.8: Covalent Bond Energies
8.10: Lewis Structures
8.11: Exceptions to the Octet Rule
8.12: Resonance
8.13: VSEPR Model and Shapes
9.1: Hybridization and the Localized Electron Model
Labs:





Molecule Polarity
Click to Run

Discussion:

re: Lewis Dot Notation
deb50duncan Nov 5, 2010 8:39 am
You will find most of your questions on the packet in chapter 7 of the new book. You will use all of the charts in this chapter. You will also need to read on ionic and covalent bonds. When drawing the Lewis Dot Structure Carbon is always in the middle, Hydrogen is never in the middle. Some of the questions in the packet you need to read carefully because the wording may be wrong. You also need to read in the book on metals and nonmetals, shared and non-shared electrons, and energy levels.
deb50duncan Nov 5, 2010 8:40 am

nchsmatts Oct 22, 2010 8:27 am
In this chapter on one of the problems that says for you to put the atoms and ions in decreasing size you really put it in increasing size. This chapter requires a lot of reading about the group and the period trends in Chapter 5. You also need to use the charts in Chapter 5 and the trends to determine the answers. Another useful tool is the periodic table in the back of the book. Some things that may slow you up are when it asks you to find which atom is more metallic, and it may give you more than one metal so you have to determine which one has more metallic qualities. You also need to pay attention to the charge above each compound when you are forming the correct Lewis Dot Notation. This chapter is just like the Lewis Dot we went over in Chemistry I. When determining the trends you will need to read all about the trends because it is not word for word in the book it takes some thought and process of elimination.

mec3539 Nov 10, 2010 8:30 am
In this lesson, make sure you read carefully to determine the correct number of electrons in a chemical. On the questions that say rank the following in terms of increasing first ionization energy; you actually rank them in decreasing first ionization energy. When looking for atomic radius, turn to page 151 in your text book. study the chart to determine the answer. Remember: When determining the largest atomic radius or ionization energy don’t necessarily look for the highest number but more study the trends. For instance Ionization energies increase across a period and decrease down a group. Be sure to read the paragraph on the question your dealing with to be sure you have the most accurate answer. when dealing with valence electrons, don’t forget that they are determined by the columns on the periodic table. Be sure not to overthink the questions. Although there are some difficult questions there are also easy ones. Don’t cheat yourself by overthinking.
Lewis Dot Notation is a fairly simple rotation contrary to popular beliefs.
Most of this rotation uses the charts in Chapter 5. You must find the trends for each chart. Most of the charts' trends increase from right to left and from top to bottom. Reading the sections "Period Trends" and "Group Trends" after each chart are very helpful.
On questions that have you put elements in order in ascending (increasing) or descending (decreasing) order, you must look at the elements relativity on the chart rather than the numbers given for each element.
One key hint is that an element is unstable if it is not octet (has 8 electrons either by itself or by sharing with another element).
This chapter also uses Chapter 6 when dealing with Lewis Dot Structures. The rotation uses some of the same problems as the Bonding and Percent rotation so if any questions may arise about Lewis Dot or something of that nature, go to the Bonding and Percent Discussion Board and read the 2 discussions posted.
-Sam Lovorn-

Dr. Gutow's Lewis Structure Tutorial

To draw a Lewis Structure follow the steps below. At each step a few molecules you can try are listed. Clicking on the highlighted molecular formulas will bring up what you should have for each molecule at each step. Try it yourself and then check as you go along. After you have mastered this there are also some quick rules which can be used for some of the atoms in the periodic table.

There are a lot of web sites with descriptions of how to draw Lewis structures. You might find their explanations or methods helpful as well. Here are a number of links in no particular order: thinkquest, Univ of Waterloo, St. Olaf College, Univ. Missouri-Rolla. You can find other links by doing a Googlesearch for "Lewis Structures".

1) Count the total number of valence electrons using the periodic table as a guide. Remember that for the main group elements the number of electrons contributed by each atom can be determined from the group the element is in. For groups 1 (1A) and 2 (2A) the number of valence electrons equals the group number. For groups 13 (3A) - 18 (8A) the number of valence electrons is the number of the group minus 10.
CH4
CF2Cl2
SO2
O3
PO43-
H2SO4
NO
2) Next draw single bonds from each of the outer atoms to the central atom in the molecule (in the examples below the central atom is in boldface type). The central atom is usually the atom with the lowest electron affinity. However, it is sometimes difficult to tell without additional information. Subtract two electrons from the total number of electrons for each bond you have made. This tells you how many electrons you have left to use elsewhere.
**C**H4
**C**F2Cl2
**S**O2
O3
PO43-
H2SO4
NO
3) Remembering that each bond counts as two electrons, put electrons on the outer atoms to give each atom a total of eight (an octet). Remember that H (hydrogen) only needs two electrons and B (boron) often only takes six electrons. Keep track of how many electrons you are using. If you run out of electrons before filling the outer atoms' octets, stop.
CH4
CF2Cl2
SO2
O3
PO43-
H2SO4
NO
4) Any electrons that were not used up in step 3 should be put on the central atom. You should now have no unused valence electrons.
CH4
CF2Cl2
SO2
O3
PO43-
H2SO4
NO
5) If any atoms do not have octets, make multiple bonds (double and triple) by sharing electron pairs from atoms that do have octets.
CH4
CF2Cl2
SO2
O3
PO43-
H2SO4
NO
6) Look for resonance structures. If you have made multiple bonds or have odd electron species where all the atoms cannot have octets, there may be more than one way to arrange the multiple bonds or place the odd electron. If so, the molecule is better modelled as an average of all the possible structures.
CH4
CF2Cl2
SO2
O3
PO43-
H2SO4
NO
7) Use formal charge to pick the resonance structure likely to contribute the most to the structure. (Formal charge) = (number of valence electrons atom starts with) - (number of electrons assigned to the atom). (Number electrons assigned to atom) = (number of non-bonding electrons) + 1/2(number of bonding electrons). Do this calculation for each atom in each resonance structure. Resonance structures with formal charges closest to zero are lower energy. Negative formal charges on the more electronegative atoms are also preferred.
CH4
CF2Cl2
SO2
O3
PO43-
H2SO4
NO
8) If the central atom is from row 3 of the periodic table or below, check for additional resonance structures where the central atom has more than 8 electrons (expanded octet). Sometimes these resonance structures produce better formal charges. This is most commonly encountered with compounds of S and P.
CH4
CF2Cl2
SO2
O3
PO43-
H2SO4
NO
Quick rules which work well for some period two atoms (H, C, N, O, F) and the halogens (group 17) most of the time.***