Where we have been:
The Quantum Model of the atom and electron configurations are a way to show where electrons are likely to be in an atom and which electrons are involved in bonding.
What we are doing now:
The Lewis Model of atomic bonding shows how the electrons are shared in Covalent Bonds. It can also be used to represent what is happening in Ionic Bonding but that is not generally done. The model includes the concept(s) of electronegativity and what these bonds are. The model is 2-dimensional but can be used to predict molecular geometry AND I will teach you how to predict the orbital hybridization even if you don’t understand what is going on by making “AXE formulas (not in this lesson though) based on lewis structures.
Miss my lesson (?!) - here’s a long video that goes over the material….https://www.youtube.com/watch?v=9VM69EgnT_o
These assignments focus on Pages. 325 -337 in the Zumdahl text.
Pay attention to my “Carney Notes”
Lesson 5 Begins Here - It covers Sections 8.10-8.12
● Watch this short review on Quick review on bond types.https://www.youtube.com/watch?v=PoQjsnQmxok
1. What determines the type of bond that is formed? Outline what characterizes an “ionic” bond vs. a “Covalent Bond”. The type of bond formed depends on the difference in electronegativity between the atoms. Covalent bond =
sharing….when the diff is less than around 1.7; ionic bond - when the difference is above about 2.1. Between these values the character is increasingly ionic and may be called ionic (esp. If a metal is involved). 2. Why do cations and anions hang out after the electron exchange?There is an
attraction between the oppositely charged ions.
3. If the difference in electronegativity is roughly 1.7 or less, then what type of bond is formed? Covalent
Carney note: He doesn’t talk about differences between 1.7 and 2.1…..fact is - they are becoming increasingly polar and may be ionic (depending on the atoms involved).
The following refer to the video re: Lewis Structures. https://www.youtube.com/watch?v=Sk7W2VgbhOg
5. How many spaces over is Carbon? 6 spaces
6. How many electrons does carbon have in it’s outer shell? (called valence electrons) The outer shell contains 2 electrons in the 2s and 2 in the 2p; a total of 4 valence electrons
7. How many valence electrons does carbon have? 4
8. What does the “octet” refer to (hint, how many electrons would make C have full shells?) The number of electrons that would make the “s” and “p” orbitals for the row an element is in full (ignoring the transition metals….these electrons go into the energy level below the row number) - the number of electrons a metal will give up OR 8-valence = number a non-metal wants to gain. The total is 8 - how many an atom wants in a lewis structure.
Look at the table he has at minute 1:10.
9. How many spaces over is B, N, O in the periodic table? 3, 5,6 10.Write the e-config for B, N, O, what’s the valence count for each?
See link below for e-configs. Count is 3, 5, 6
https://www.google.com/search?q=electron+configurations+of+all+elements&oq= electron+con&aqs=chrome.2.69i57j0j69i59l2j0j69i61l2j69i60.2859j0j7&sourceid= chrome&ie=UTF-8
11.How many dots are drawn around B, N, O atoms? 3, 5, 6
Carney note: the transition elements are decidedly absent here…...metals don’t generally form covalent bonds AND the electrons of transition metals are mostly in the “d” block….which has more than an octet possible! The Lewis Model describes
covalent bonds which are generally between non-metals.
12.Describe the pattern from the above three questions. The number of valence electrons = the dots
13.How many electrons are in a covalent bond as represented by a dash? Two electrons are in each bond
14.Why does Nitrogen make three bonds and have one “lone pair”? It has 5 valence electrons - that means it needs 3 more from shared atoms - so 3 + 3 = 6; two are left that it has already in the “lone pair”.
15.How many bonds does fluorine make? How many lone pairs does F have? 1, 3 16.A dash represents a single bond and is called a ____ bond? sigma
18.If there is a triple bond, state the quantity and type of bonds between them. 1 sigma and 2 pi bonds
Skip the “Formal charge stuff from 2;47 to 3:24 in this video - it’s not great and I have MUCH better ones or you later.
19.Elements in the second row - C, N, O, and F want 8 electrons around them. 20.Hydrogen can’t have more than 2 total.
21.Name two elements that can have more than 8 electrons. Any element in the p block beyond the second row…..Si, P for example.
22.Check your comprehension and comment on how you did. This varies by student of course.
Steps to write Lewis Structures
Watch this video and answer the questions: https://www.youtube.com/watch?v=1ZlnzyHahvo
23.Write out step 1 determine valence electrons
24.How many did hydrogen have; how many did oxygen? 1, 6
25.For CN-, there were 10 electrons - why?c has 4 and N has 5 but it’s negative so an extra one there - that sums to 10
26.Write out step 2. Total valance electrons then put 8 around each atom (except H)
27.Which atom goes in the center of the structure? Usually it’s the first in the compound; the one with more bonding locales.
28.Where does H ALWAYS go? On the outside
29.Write out step 3 (see Below and the Bonehead method written out after these questions)
30.Write out step 4
Steps 3 and 4 above and 5 below involve making sure they all (except H that gets 2) have 8; if they don’t, re-arrange by removing two from each of two atoms and then putting in center between these same atoms to make a double bond. I just start by giving ALL atoms 8 (or 2 for H) then counting. If too many - reduce by making double or triple bonds. Keep in mind electronegativities of atoms and who might get the double bond preferentially.
31.Why is the H outer shell filled with only 2 electrons (not in video - think back to the last lesson). It’s electrons are in the 1s energy level - and that only has a sphere (s) so max is 2.
32.Write out step 5 and when it is necessary. (see above)
34.How many sigma, and pi, bonds does O2 have? 1 sigma and 1 pi.
Carney’s “Bonehead” method of doing the structures: I’ll show you a slightly different way in zoom I call it “the bonehead method” ….here’s my steps. 1: add up all the valence electrons; 2) Choose the central atom (It’s usually the first one in the compound unless that is H, else, put the more electronegative one there and if forced to because there are multiple seemingly equally plausible options, calculate formal charges - see below later ); 3) give H two and all others 8 around them; 4) add up the total dots, if it equals the count from (1) you are done BUT if it is over eliminate two on each of two atoms and make a double bond between them then recount….repeat if needed until the valence count equals the dot count. Don’t forget….If an ion, you must put brackets around the structure and put the charge of the ion on the outside as a superscript. I’ll model this in my
lesson/lecture on the subject.
Carney Note on exceptions to the octet rule: Elements in the 3rd row and beyond like P, S can have more than 8 electrons because their electrons can go into the next d orbital. Example, P and S can put electrons in the empty 3d orbital. Ga, and Br could put electrons in the empty 4d orbital. They have these d orbitals in this energy level available…..electrons are already occupying the s and p sublevels….they just don’t normally put electrons there. BUT, they could go there and they do in some
compounds. Others have odd numbers (like N) an H doesn’t want an exception….what about these.
Watch this video that explains this point and does a good job explaining what’s going on between atoms:
https://www.youtube.com/watch?v=Dkj-SMBLQzM
35. Write a few notes below. Varies by student
Carney on Resonance Structures:
Sometimes, to balance the bonds to meet octets, there are multiple places to put a double bond but you have to choose one. See page 333 in your text - you’ll see the double bond between the O and N can be seen to rotate. This is called “resonance”.
● Put the individual structures in brackets. ● Put arrows between them.
● It’s a model - note - resonance doesn’t really happen like the model
Formal Charge and which structure is best
Watch this video on how to calculate Formal Charges. https://www.youtube.com/watch?v=vOFAPlq4y_k
Formal charge can be used to assess possible structures for likelihood.
In the event you can make more than one structure - like, which goes in the middle of the SCN- molecule (it’s an ion)? The S, the C, or the N. Now watch this video on using
Formal charge to evaluate which is better.
https://www.youtube.com/watch?v=Dj4sz6L9IdM
You can make three structures; He’s got S-C---N; S=C=N; there is another N-C---S…..I’ll do in Zoom I think along with these others.
- Note: add up electrons on each atom where dots count as 1; dashes count as ½. - The formal charge on an atom is = (valence # - sum of electrons)
- The sum of all charges on atoms must add up to the charge of the molecule (zero or the charge on the ion).
- The best model is the one: 1) with the most zero formal charges; 2) the more electronegative atoms get the negative formal charge.
Now read sections 8.10 - 8.12 in your textbook and do these problems in your textbook: Page 353d - e: 85, 87, 89, 91, 93, 104, 105a-d
All odd solutions are in your textbook 104:
https://www.google.com/search?q=lewis+structure+pocl3&tbm=isch&source=iu& ictx=1&fir=GeUeDfbUVxWnKM%253A%252Cesqy11VcrgIyQM%252C_&vet=1&usg =AI4_-kTh2zmBnIr0nhi1EcmHYGuD7ZdqXw&sa=X&ved=2ahUKEwj644DVrfroAhW R4J4KHXX6DXAQ9QEwBnoECAUQLA#imgrc=GeUeDfbUVxWnKM:
B. See pages 335-336 in your text where they solve it. Note that the best one will end up with double bonds - then you have resonance structures.
C. - Play the same game with the perchlorate ion as the sulfate ion - you will end up with formal charge favoring one with some double bonds.
D.Play the same gave for the phosphate ion as for the sulfate ion…..and you wlil end up with a double bond on one oxygen; the other will have a formal charge of -1. This has more with fc of 0.
VSEPR Theory and The geometric (3D) shape of molecules. You MUST be able to do Lewis Structures to do this. Also, we will return to the Quantum Model and talk about how the electrons are shared in Covalent Bonds (in what are called Hybrid
Orbitals….to understand it you MUST be able to do orbital diagrams and make electron configurations well) BUT you can predict them even if you don’t understand what’s going if you can do Lewis Structures and predict shapes. Each shape is associated with a particular orbital hybridization you see…..(Truthfully, this last subject is quite abstract and my goal is more that all of you can predict and about 50% understand the first time around in my class….)
Miss my lecture - then watch these.