8.2 Lewis Dot Structures Flashcards by Tyler Winn

Lewis Dot structures Way of representing the valence electrons
* doesnt even look at the core ones
* remember, the valence electrons are the ones in charge of chemical reactions and the making and breaking of the bonds

This lession is primarily going to be for covalent bonding and not ionci

Group 1 = 1 valence electron
Group 2 = 2 valence electrons
Group 13 = 3 valence electrons
Group 14 = 4 valence electrons
Group 15 = 5 valence electrons
Group 16 = 6 valence electrons
Group 17 = 7 valence electrons
Group 18 = 8 valence electrons (except He which just has 2 valence electrons)

How well did you know this?

Not at all

Perfectly

Atoms try to organize themselves in a way where they fill that octet (thats the octet rule)

atoms will typically transfer or share electrons in an attempt to have 8 electrons around them.

In an ionic compound atoms are going to transfer electrons to complete the octets

In covalent compounds atoms are going to share electrons to complete octets

How well did you know this?

Not at all

Perfectly

So when you have a metal with a low ionization energy (doesnt take much to pull that electron off) next to a non metal w/ a high (very negative) electron affininity = the transfer of an electron = ionic compound
* again, its very negative because when it steals can electron it releases energy from the system in an exothermic reaction - since the system is losing energy its considered negative

So below shows Cl stealing Na’s electron to fill its own octet and leave Na w/ and empty valence shell. So they’re both happy. And now they hangout because one is positive and the other is negative, meaning they form an ionic compound
* Keep in mind, they arent sharing electrons here, one is just positive because it lost an electron and the other is negative because it gained an electron

How well did you know this?

Not at all

Perfectly

The below is showing 2 Cl atoms that are going to be covalently bonded
* the problem is that both are missing 1 electron to fill its octet, so the must share 1 each to fill the octet

Both Cl’s have a realtively high ionization E, meaning neither of them want to lose an electron. They have the same electronegativity meaning neither will lose an electron

The ones that are shared are called bonding electrons

The ones that are not being shared are called non bonding electrons
* also called lone pairs
* Each cl has 3 lone pairs (or non bonding electrons)
* typically bonding electrons are lower in E than bonding <— this is the driving force for bonding in the first place - its an exothermic process for making a bond (the system loses energy, meaning its more stable) and its an endothermic process to break a bond (meaning you need E from the surroundings to break an already formed bond, adding E to the system)
* So its energetically favorable to make bonds

How well did you know this?

Not at all

Perfectly

Major exceptions to the octet rule
* He = 2 valence electrons (doesnt want filled octet)
* This makes tons of sense, because its only 1s^2 at He
* Hydrogen only holds 2 valence electrons (not 8) because it wants to look like its nearest noble gas (which is He)
* Al and Be can be involved in covalent bonding (even though they’re metals) - will be polar covalent bonding
* Be has 2 valence electrons - typically will share 1 with one atom and share another w/ another atom and make 2 bonds. So Be can only get 4 electrons around him, cannot get a filled octet
* B/Al only have 3 valence electrons, so can only make 3 bonds, meaning the max electrons it can get are 6. However, these 2 can be found w/ a filled octet, its just uncommon

We can actually predict how many bonds an atom is going to make. It typically depends on how many electron short of an octet they are
* For instance, carbon has 4 valence electrons, so it will form 4 bonds to complete its octet

How well did you know this?

Not at all

Perfectly

Another exception to the octet rule is an expanded octet
* a typical shell is going to have an s orbital and 3 p orbitals (8 total)
* however, there arent just s and p orbitals but also d and f orbitals
* but you only start getting those d orbitals and shell number 3 and f in shell # 4
* that means only elements in period 3 or lower can start throwing electrons in the D orbital, giving them an expanded octet
* So starting at Na and over / down can potentially have this expanded octet. However, they often dont.
* So sulfur often makes 2 bonds, following the octet rule, just like oxygen. However, its not uncommon to see sulfur making 6 bonds, obviously breaking the octet rule. This is because it can fill that d orbital. (that would be 12 electrons)
* Oxygen just doesnt have this capability

How well did you know this?

Not at all

Perfectly

Another way to break the octet rule is simply having an odd # of electrons

So the molecular compound NO has 11ve-, meaning they cant all be paired

this is fairly uncommon

You’ll typically give more electrons to the more electronegative atom first, which in this case is oxygen which is why it filled its octet
* remember this

An atom that is more elctronegative will release more energy when it steals an electron creating an exothermic reaction, meaning the system gets negative energy which is why its called electronegative
* also electronegativity refers to how strongly an atom attracts electrons in a bond, not just how much energy is released when it gains an electron on its own (which is electron affininty)

How well did you know this?

Not at all

Perfectly

Halogens have 7ve-, meaning they’re only 1 short of having a filled octet. Meaning we can easily predict how many bonds they’ll make (1) to entirely fill their octet

How well did you know this?

Not at all

Perfectly

The atom in the middle generally makes the most bonds

The least electronegative element generally can make the most bonds, and therefore would go in the middle of the molecule
* Remember, F is the most electornegative and can make 1 bond
* Carbon is significantly less electronegative and can make 4
* Hydrogen is an exception, it has low electronegativity but can only make 1 bond

If you’ve got more than 1 halogen in a molecule, well the lower ones are less electronegative (further from florine) and will make more bonds - so will likely be in the middle

How well did you know this?

Not at all

Perfectly

Drawing lewis structures
1) Set up skeleton w/ single bonds (central atom is the atom which can make the most bonds)
2) Fill the octets of the outside atoms with lone pairs of electrons - do this first
3) Any remaining valence electrons go on central atoms (since we fil the peripheral ones first this one will often not be full, meaning we wil have to adjust bonds)
4) Once all electrons are used, form multiple bonds to central atom if not full

So once you fill up those outside atoms, ask yourself if you have any extra Ve

How well did you know this?

Not at all

Perfectly

Draw the lewis structure for NF3
* typically the first element (nitrogen in this case) is the most electronegative element. Meaning it likely will be the central element because less electronegativity means it can form more bonds (further from F = decrease Ve, meaning it will want to grab more Ve to make more bonds)

Start by filling up the outside atoms (or filling your more electronegative atoms first, because those are going to be the outside ones)
* rememebr, electronegativity has to do with how much they like to pull electrons towards them, so it would make since that the more electornegative atoms have the full valence shells

So you put any extra electrons on that central atom. Then ask yourself, is that central atom happy/full. In the case below nitrogen follows the octet rule so its happy.

How well did you know this?

Not at all

Perfectly

Draw the lewis structure for HCN
* so for nonbinary (more than 2 elements) molecules, they typically actually put the one in the middle thats supposed to be in the middle
* So in this case its carbon

How well did you know this?

Not at all

Perfectly

Hydrogen has 1 electron in nature, but will bond to get 2 electrons to match its nearest noble gas (HE)

As an ion hydrogen loses its 1 electron so it has no bonds and forms a +1 charge. So its not actually matching He’s configuration because helum has 2 electrons and at this state hydrogen has just a proton (which is why hydrogen is often just called a proton)

How well did you know this?

Not at all

Perfectly

Draw the lewis structure for CO2

carbon is less electronegative, meaning it can make more bonds, so well put it in the middle

So i started by filling the peripheral atoms octets first. I saw that the central wouldnt have enough so will have to start making double bonds
* Found that oxygen likes to double bond and carbon likes to have 4 bonds

I didnt want to do a triple bond w/ Oxygen because oxygen likes to create 2 bonds

How well did you know this?

Not at all

Perfectly

What is it called when molecule can have multiple structures drawn for its fomrula?

Resonance
* when you have resonance you have delocalized electrons.
* These are electrons that are in more than 1 location at the same time
* so below the real structure is some average of all the structures, but some of the structures contribute more than others
* To figure out which ones contribute more you’re going to worry about formal charge
* when you’ve got different resonance structures, the best one has the least formal charge

How well did you know this?

Not at all

Perfectly

Formal charge formula: Number of valence electrons for an atom (on the periodic table) - [1/2 bonding e- + nonbonding e-]
* shown below

You can also do #of Ve on the periodoic table - # of dots and lines

So for the example below it would be 6 valene electrons minus 6 dots + 1 line = -1

Draw the lewis structure for N2O

So this is tricky because you’d want to make it symmetrical and put oxygen in the middle. However, the less electronegative element actaully goes in the middle which is nitrogen
* also, remember, the one thats written first in a binary compound often goes in the middle

Formal charge = Ve listed on periodic table - dots + lines (lines only count as 1)
* formal charges will always add up to wahtever charge your compound has. In this case its neutral

So there was 3 possible resonance structures for this

So formal charge rules listed on next slide.

Were trying to figure out which ressance structure is better between the top 2. The bottom was ruled out because it had a high number formal charge in it (+2)

If the first 2 rules dont work follow the last one which states that the negative charge should be on the more electronegative atom (this makes since because the more electornegative atom is the one that would grab the electron, thereby making it more negative)

So the highlighted one below is the major resoance contributer
* however, when you have ressoance structures, none of them are perfectly correct. The molecule would really look like some average of all of the ressoance structures. However, it would look more like the major contributor (highlighted below) than it would look like the other 2.
* Delocalization = electrons are in multiple locations at the same time. So they’re partially in multiple locations at the same time.

Formal charge rules:
* Formal charge will = whatever charge is on compound - in neutral compounds it will be neutral
* 1) Fewer formal charges on atoms.
* 2) Lower magnitude formal chaerges (closer to 0, 1 is better than 2). -1 better than -2
* 3) Negative charge on more electronegative atom.

follow these rules in order

Draw the lewis structure for SF4

The least electronegative element goes in the middle (because the more electronegtaive ones just want to grab those electrons)
* also less electronegative ones typically can form more bonds (because they’ll have a less full octet)

So sulfur is less electronegative so will go in the middle

We after filling the peripheral atoms we find that everyones octet is full w/ 2 left over.

However, sulfur is in the third row (meaning it has a d orbital), so it can go over the octet rule
* We will find that the only time it has a chance of going over the octet rule is when its the central atom and 3rd row or lower - however it doesnt have to go over the octet rule, but it sometimes will. in this case it does

So sulfur has 10 electrons to use all the valence electrons

At this point you should check your formal charges, because you might be able to make it lower by forming a double bond instead of adding that lone pair onto S
* this will only be the case if you have a negative formal charge atom on the outside, bonded to the central atom w/ a positive formal charge
* however we have no formal charges below

Draw the lewis structure for XeF4
* this is weird because its got a noble gas in it.
* remember, we said that the only time nobles gases bond is when they’re bonding to something very electronegative (something that really really wants an electron), like F, but even then it doesnt haoppen very often

If you have a noble gas in your compound its going to be the central atom - because it will be less electronegative than anything else, because its got a full octet and isnt looking for electrons

If you’ve got a noble gas, thats got a full octet, and going to make bonds, its going to exceed the octet rule.
* remember, its only that central atom that exceeds the octet rule, and F isnt even a period 3 element or lower

No formal charges on the F or Xe

Draw SO4^2-

So this is not the greastest structure in the world, w/ all the different formal charges.
* however, if you’ve got an atom that can have an expanded octet in the middle, and its got a positive charge of some sort, and the ones next to it on the outside have a negative formal charge. Well you can actaully reduce your overall formal charge by having your negatives donate electrons, and form more bonds. So were going to have the negatives share their electrons w/ the positive to even it out.
* So again, this only worked because sulfur can break the octet rule, and we used this to reduce its overall formal charge. S now has 12 electrons
* NOTE: if we donated anymore electrons to S from the other 2 it would give S a negative formal charge, and oxygen is more electronegative than F so we would rather have that formal charge on the O than the F (one of our rules)

Also when you have an ion put the entire lewis structure in brackets and put its charge

Draw the lewis structure for NO3^1-

Nitrogen is less electronegative so it goes in the middle

So below theres not 1 major resonance contributor because they all have the exact same formal charges, even though theres 3 possible ressonance structures

resonance always implies delocalized electrons
* none of the 3 resonance structures below is an accurate dipiction of reality - none really exist. The real molecule is some average of all 3 at the same time. Its not switching back and forth between these 3, its litteraly all of them at the same time

NOTE: double bonds are stronger and shorter than single bonds (they pull the atoms together more)
* however, when we look at the sturcture of NO3^1- we find that all 3 bonds are the exact same length all the time. Its not like its flipping back and forth between all 3.
* So its not like any of the structures drawn below give an accurate dipiction of what the molecule looks like at any 1 moment in time.
* the molecule looks like the average of all these structures
* so you can think of this as a partial double bond between all of them (shown at the bottom). Its not as strong or as short as a double bond between any of them, but its shorter and stronger than a single bond between all of them. its like a hybrid
* its like 4 bonds spread across 3 locations, its really like 1 and 1/3 bonds.

So the Nitrogen has a fully positive partial charge because all 3 of the ressoance structures have it as a + charge

ech oxygen has -2/3 of a charge (represented w/ a partial charge symbol below on the resonance hybrid)

The bonds below being 1 1/3 is what we mean by delocalized electrons. They’re in multiple different places at the same time.

Expanded octet seems to really only apply to P block elements row 3 and lower (basically starts w/ Al)
* however, they must be a central atom
* The reason its the P block elements is theoretically they can use low-energy d-orbitals to accommodate more than 8 electrons

Even though Al and Si don’t have electrons in the 3d orbitals normally, the 3d orbitals still exist, and can be used to make room for more electrons when forming certain compounds

For Br the ground state is [Ar] 4s^2, 3d^10, 4p^5.
* when it wants to add extra electrons it can use the 4d orbitals

So for Al its electron config is [Ne] 3s^2, 3p^1, however it can use the higher energy 3d orbitals to acccomdate extra electrons when it exceeds the octet rule

hydrogen does not ever become the center atom. so doens tfollow our electronegativity rules

Remember, typically the center atom is the least electronegative atom
* Electronegativity is about how strongly an atom pulls electrons toward itself in a bond
* An atom w/ high electronegatvity tends to attract electrons strongly in bonds, meaning its pulling electrons toward itself, often resulting in polar covalent or even ionic bonds
* Electronegative atoms don’t like sharing and central atoms like to share - electronegative atoms prefer to “hog” electrons, not donate or share them

So in our case Si is considered central because H cannot be because it only makes 1 bond.

All the below have neutral partial charges

the answer is A

Hydrogen can’t go in the middle so Al must
* also, in binary compounds the first one is typically the central atom

Al doesnt follow the octet rule beacause it only has 3 valence electrons, so can only pair 3 times
* only has 3 electrons, so it typically only forms 3 bonds, meaning it only gets 6 electrons instead of 8
* its also not very electronegative, so it doesnt strong attract electrons
* Itsualy fine w/ 6 electrons -