Lecture 2b - Molecular Dot Structures

Summary

This lecture delves into molecular dot structures, essential for understanding molecular bonding and reactivity. It outlines a step-by-step process to draw dot structures, emphasizing the importance of order. The lecture explains how to calculate valence electrons, determine the central atom, and distribute electrons to achieve octets. It introduces the concept of resonance, where electrons are delocalized, and the use of formal charge to predict molecular structure. Exceptions to the octet rule and expanded octets are also discussed, providing a comprehensive foundation for predicting chemical behavior.

Takeaways

• 🔬 Dot structures for molecules are crucial for understanding bonding and predicting reactivity and chemical behavior.
• 🧩 The process of drawing dot structures involves a series of steps that must be followed in a specific order to ensure accuracy.
• 💡 The first step in creating a dot structure is calculating the total number of valence electrons available from all atoms in the molecule.
• 🚫 Hydrogen is an exception in dot structures; it only needs two electrons to achieve a stable configuration, similar to helium.
• 🔑 The central atom in a molecule is typically the first non-hydrogen atom listed in the chemical formula, around which other atoms are bonded.
• ✅ Single bonds are used initially to connect atoms to the central atom, with each bond representing two electrons.
• 🔄 After initial bonding, lone pairs are added to outer atoms first to complete their octets before considering the central atom.
• 🔄 If there are leftover electrons after outer atoms' octets are filled, they are added as pairs to the central atom.
• 🔁 Resonance structures occur when there is more than one valid way to distribute electrons, especially in molecules with delocalized bonding.
• ⚖️ Formal charge calculations help determine the most stable resonance structure by identifying which atom can best handle a multiple bond.
• 💥 Expanded octets are exceptions to the octet rule, where certain elements, like phosphorus and sulfur, can have more than eight electrons around them.

Q & A

What is the primary purpose of creating dot structures for molecules?

Dot structures for molecules are primarily used to determine the bonding present in a molecule, which in turn helps in predicting its reactivity, where reactions will occur, and how they will happen.

How does shape prediction in chemistry relate to molecular dot structures?

Shape prediction in chemistry is related to molecular dot structures because the shape of molecules, influenced by their bonding patterns, is crucial for understanding chemical behavior, such as drug interactions within the body, which are based on the 'lock and key' model of molecular shapes.

What is the significance of the order in which the steps for drawing dot structures are performed?

The order of the steps for drawing dot structures is significant because changing the order can lead to incorrect dot structures and incorrect predictions about the molecule's bonding and reactivity.

Why is the total number of valence electrons calculated first when drawing dot structures?

The total number of valence electrons is calculated first to determine the total number of electrons available to fill the octets of all the atoms in the molecule.

Why is hydrogen treated differently when considering valence electrons in covalent bonds?

Hydrogen is treated differently because when it forms a covalent bond, it is considered to have two electrons around it, resembling a noble gas configuration, and thus it is satisfied with just one bond rather than needing to achieve an octet.

What is the general rule for identifying the central atom in a molecule for dot structure drawing?

The central atom in a molecule for dot structure drawing is generally the first non-hydrogen atom listed in the chemical formula when read from left to right.

How are lone pairs added to atoms in a dot structure, and in what order?

Lone pairs are added to complete the octets of the outer atoms first, followed by the central atom if there are electrons remaining. The order is important to ensure that all atoms achieve a stable electron configuration.

What is resonance in the context of molecular dot structures, and why is it important?

Resonance refers to the phenomenon where a molecule can be represented by more than one valid dot structure due to delocalized electrons. It is important because it reflects the actual bonding situation in a molecule more accurately than a single dot structure and influences the molecule's stability and reactivity.

How does formal charge help in determining preferred resonance structures?

Formal charge helps in determining preferred resonance structures by indicating which atom can best handle a multiple bond, with the preferred structure having formal charges closest to zero for all atoms.

What is an expanded octet, and which elements are known for forming them?

An expanded octet is a situation where an atom has more than eight electrons in its valence shell, typically seen with non-metals later in the periodic table, such as phosphorus and sulfur.

Why might the sum of formal charges not equal the overall charge on a molecule if calculated incorrectly?

If the sum of formal charges does not equal the overall charge on a molecule, it indicates an error in the calculation of formal charges, which can lead to incorrect predictions about the molecule's structure and properties.

Outlines

00:00

🔬 Introduction to Molecular Dot Structures

This segment introduces the concept of molecular dot structures, emphasizing their utility in determining molecular bonding and reactivity. It outlines the significance of dot structures in predicting chemical behavior, particularly in medicinal chemistry where the 'lock and key' mechanism of drug interactions is highlighted. The lecture sets the stage for a detailed exploration of the steps involved in drawing dot structures for molecules, noting the complexity of molecules over atoms and the importance of following a systematic approach.

05:01

🧠 Steps for Drawing Molecular Dot Structures

The speaker elaborates on the systematic steps required to draw molecular dot structures. The process begins with calculating the total number of valence electrons from all constituent atoms, with a special consideration for hydrogen. The central atom, typically the first non-hydrogen atom in the chemical formula, is identified and connected to other atoms with single bonds. The speaker stresses the importance of keeping track of electrons used in forming these bonds and subtracting them from the total valence electrons. The subsequent steps involve adding lone pairs to outer atoms to complete their octets and, if electrons remain, adding them to the central atom. The lecture also touches on the possibility of forming multiple bonds if necessary.

10:01

🔍 Examples and Resonance in Molecular Structures

This part of the lecture delves into practical examples such as methane (CH4), water (H2O), and others, with the promise of a separate video detailing the dot structures for these molecules. The concept of resonance is introduced, explaining how electrons can be delocalized over multiple atoms, leading to different resonance structures. The speaker discusses how resonance structures can be represented and how they contribute to the overall stability of a molecule by lowering its energy. The segment also hints at the use of formal charge calculations to determine preferred resonance structures.

15:03

📐 Formal Charge and Exceptions in Dot Structures

The final segment discusses the calculation of formal charge to determine the most stable resonance structures. The process involves assigning electrons from bonds and lone pairs to atoms and calculating the formal charge by subtracting the assigned electrons from the valence electrons of the atom. The preferred structure is indicated by formal charges closest to zero. The speaker also mentions exceptions to the octet rule, such as beryllium preferring a tetravalent structure and boron a hexavalent one. The lecture concludes with a mention of expanded octets and odd-numbered electron species, particularly relevant in fields like atmospheric chemistry, and encourages viewers to watch example videos for further clarification.

Keywords

💡Dot Structures

Dot structures, also known as Lewis structures, are diagrams that represent the valence electrons of atoms in a molecule. They are used to predict the bonding and reactivity of molecules. In the script, dot structures are central to understanding how molecules form bonds and how they might react, as they help visualize the distribution of electrons and the potential for chemical reactions.

💡Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom, which are involved in chemical bonding. The script emphasizes the importance of counting valence electrons to determine the total number available for bonding in a molecule. This is crucial for constructing accurate dot structures and predicting the molecule's reactivity.

💡Octet Rule

The octet rule states that atoms tend to combine in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. The script mentions that atoms aim to achieve an octet through bonding, which is a key principle in understanding how molecules form and the resulting dot structures.

💡Central Atom

In a molecule, the central atom is the atom to which other atoms are bonded. The script explains that when drawing dot structures, the central atom is usually the first non-hydrogen atom listed in the chemical formula. This concept is important for determining the connectivity and structure of the molecule.

💡Lone Pairs

Lone pairs refer to a pair of electrons that are not involved in bonding and are instead located on an atom. The script discusses how lone pairs are added to outer atoms first to complete their octets. Lone pairs are significant in dot structures as they represent the non-bonding electrons that can influence a molecule's reactivity and shape.

💡Formal Charge

Formal charge is a way to keep track of the charge distribution in a molecule by assigning electrons from bonds to individual atoms. The script introduces formal charge as a tool to determine the preferred resonance structure in molecules where delocalized electrons exist. It is calculated by considering the number of valence electrons an atom would have, minus the electrons assigned to it in the molecule.

💡Resonance

Resonance occurs when a molecule can be represented by more than one valid Lewis structure, with electrons being delocalized over multiple atoms. The script uses resonance to explain situations where electrons are not localized between two atoms but can be distributed among several. This concept is important for understanding the stability and reactivity of certain molecules.

💡Multiple Bonds

Multiple bonds, such as double or triple bonds, occur when two or more pairs of electrons are shared between two atoms. The script mentions that if there are not enough electrons to complete the octets of all atoms, multiple bonds may form. These bonds are significant in dot structures as they alter the molecule's reactivity and shape.

💡Expanded Octets

Expanded octets refer to atoms that have more than eight electrons in their valence shell, which is an exception to the octet rule. The script notes that some elements, like phosphorus and sulfur, can have expanded octets. This concept is important for understanding the bonding in molecules that do not follow the typical octet rule.

💡Odd Numbered Electron Species

Odd numbered electron species are molecules that have an odd number of valence electrons, which is less common but can occur in certain chemical environments. The script briefly mentions these species in the context of atmospheric chemistry, where high-energy radiation can lead to the formation of such species. Understanding these species is important for studying chemical reactions in extreme conditions.

Highlights

Introduction to molecular dot structures as an expansion from atomic dot structures.

Dot structures are essential for determining bonding in molecules, which in turn helps predict reactivity and reaction outcomes.

Dot structures can also assist in shape prediction, crucial for fields like medicinal chemistry where drug interactions depend on molecular shape.

A step-by-step process is outlined for drawing dot structures for molecules, emphasizing the importance of following the steps in order.

The first step in the process is to calculate the total number of valence electrons available for bonding.

Hydrogen atoms are an exception, needing only two electrons to achieve a stable configuration.

The central atom in a molecule is typically the first non-hydrogen atom listed in the chemical formula.

A single bond is used to attach each atom to the central atom, and the number of electrons used in these bonds is tracked.

Lone pairs are added to outer atoms first to complete their octets, with any remaining electrons added to the central atom.

If there are not enough electrons to complete the central atom's octet, multiple bonds may be formed with outer atoms.

Examples of dot structures for methane (CH4), water (H2O), and other molecules will be provided in a separate video.

Resonance structures are introduced as a concept where electrons are delocalized and not confined to a single bond.

Resonance structures can lower the energy of a molecule and are represented by dashed lines indicating partial bonds.

Formal charge is used to determine which atom can best handle a multiple bond and to predict the preferred structure among resonance forms.

The formal charge calculation involves subtracting the number of assigned electrons from the valence electrons of an atom.

The sum of all formal charges in a molecule should equal the overall charge of the molecule.

Exceptions to the octet rule are discussed, such as beryllium preferring four electrons and boron preferring six.

Expanded octets are mentioned, where atoms like phosphorus and sulfur can have more than eight electrons due to additional bonds.

Odd-numbered electron species are briefly discussed, which can occur in high-energy environments like the stratosphere.

The lecture concludes with a reminder to watch example videos for a better understanding of the concepts discussed.

Transcripts

welcome to the second part of uh lecture

two

uh so to be we are going to talk about

uh molecular dot structures so we've

talked about

how to do atomic dot structures and now

we're going to look at expanding that a

little bit

and moving on to how do we do dot

structures for entire molecules

so when we're doing dot structures uh

for molecules uh we find

reasons we want to do them is because

the dot structures themselves are very

useful

for helping to determine the bonding

that is present in a molecule and if we

can determine the bonding that's present

we have a good chance of determining its

reactivity

where reactions will happen how

reactions will happen where will things

bond

through a reaction and so this is really

a very powerful tool a lot of times it's

the starting point

in determining what kind of products

you're going to get out for for

different reactions

the other thing it can help us do is to

lead to shape prediction

and shape prediction in chemistry is is

very important

if you are in something like a medicinal

chemistry field

uh if you're familiar with uh with the

lock and key idea of how

how drug interaction occurs within the

body all of that is due to

shape of the ends of molecules or the

middles of certain

uh certain molecules and yeah we'll see

some ideas of how how that comes about

so shape prediction gives us an idea of

chemical behavior

like we just talked about so there's a

series of steps that we need to use in

order to draw

dot structures for molecules not quite

as easy as what we did for the atoms

and that makes sense molecules more

complex than atoms so the things that we

do

with molecules tend to be more complex

than what we would do for atoms and so

we're going to look at this

series of steps and the important thing

with this series

of steps is that we do them in order if

you do them

in a different order you say well just

do these steps in different order

you're not going to get the same answer

and your dot structure will most likely

not be correct so the first thing that

we're going to do is we're going to

calculate just a total number of valence

electrons

now the reason we do that is we have a

collection of atoms that we're going to

make a molecule

out of and each of them has a certain

number of valence electrons

they all want to try to achieve their

octet to have those

eight valence electrons count around

them that's why they form ions when

they're doing ions but here

when we're doing dot structures

generally we're going to use them for

covalently bonded atoms and so what

we're going to do is we're going to take

all these valence electrons we're just

going to put them into a big pot

and say okay these are the electrons we

have available

we have to fill all of the octets of all

the atoms with this

number of electrons now the one

exception to that of course is

the hydrogen atom anytime you attach a

hydrogen atom with a bond

it feels like it has two electrons in it

because each of those covalent bonds

represents two electrons

and so uh in that case we just have

two electrons around hydrogen that makes

it look like helium

it feels like a noble gas and it's happy

so hydrogen once it makes its bond

uh it's perfectly happy okay so we've

added up all of our valence electrons we

want to write that number down

uh the next thing we want to do is we

want to draw the molecule

attaching each atom to a single central

atom

uh with a single bond now you may ask

what is the central atom going to be

generally

it's the first non-hydrogen atom listed

in the formula okay so when you look at

a chemical formula

if it's written in a kind of an accepted

manner

the first as you read left to right the

first atom that is not hydrogen

is going to be the atom that is in the

center

so we put that atom in the center we

attach

all of the other atoms to it with a

single a single bond

now the other thing that i didn't put on

here is you do have to keep a count

of all of the electrons that you've used

and every time that you add a bond

you've used two electrons and so when

you attach

all of those other atoms to that central

atom you have to count the number of

single bonds that you've put on there

and uh multiply by two to get the number

of electrons you used to make those

single bonds and you want to subtract

those

from the original total number of

valence electrons

okay so now you've got a certain number

of electrons uh

that you have left to try to complete

everybody's octets in there so the first

octets that we're going to

fill are the outer atoms so you have the

inner atom the one in the center

and then you have all of the outer atoms

so we're going to add lone pairs to the

outer atoms first

we're going to complete their octets now

it's very important we're going to

complete their octets

we don't worry about how many valence

electrons does the lone atom have

okay because lone atoms have a certain

number of electrons around them but when

they form a compound they do so to fill

their octets

so we're going to add lone pairs those

outer atoms to complete their octets

if after you've finished the octets of

the outer atoms and you still have

electrons left over

you will add them as pairs to the

central atom and this is very important

whenever you're adding electrons here

you add them as pairs you don't add

singles in different places you just add

them

as uh pairs okay now

if you put all your electrons around

the outer atoms and you have no

electrons left which

happens a lot of times sometimes you

will find

that your central atom does not have an

octet and if that happens

then you have to start forming multiple

bonds

with one of the outer atoms or more of

the outer atoms or it may need

form a double bond or a triple bond or

whatever it takes we'll look at some

examples to see how to do that

okay so here's some examples that we're

going to go through

uh methane ch4 water h2o

uh cobr2 and clf2 minus

uh i will go through those in a separate

video

i will post the link uh to that along

with the link to this video as well so

you will see

uh some example videos with dot

structures

another example of something we'll look

at is something called resonance

when we're looking at resonance and we

look at the examples that we just did

on in our link there we may assume that

all electrons that get shared between

two atoms are

localized and when i say localized that

means that those

electrons remain only between those two

atoms

and this isn't always the case okay this

isn't always the case

so what can happen uh is that you have

uh

molecules where uh we have different

options

on where to place other pairs of

electrons so

if we reach a place in our dot structure

where we have a choice onto where to put

a double bond we can put it between this

atom this atom or this atom in this atom

or this atom and this atom

we have a situation where we have uh

something called resonance okay another

term for resonance is

delocalized bonding and delocalized

means that

a pair of electrons does not have to

exist just between

two atoms that can exist between many

different pairs of atoms

so when we're looking at an actual

structure there's lots of ways to look

at it we can look at it as

a compilation of all of the different uh

resonance structures which sometimes

you'll see

or we can look at it in terms of a

hybrid

of all of the various structures and any

time you have delocalization of the

bonds it lowers the energy of the

molecule

and so that's why a combination of all

three of

the structures is energetically more

preferred

uh than just three therefore however

many different

uh individual uh structures you have so

here's an example

of uh of ozone uh so we've got a

structure on the left

there where we had a double bond

over here on the left and on the right

hand structure we put the

double bond over here on the right okay

and which one is correct well they're

both correct they're actually

both equivalent and so if we're looking

at resonance we can

show that we have two different forms

like that

that's one way of showing all of the

resonance forms

or you can also show it with a dashed

line

you can kind of see right here and right

here

a dashed line that indicates that that

bond is

actually a partial bond between

all three of those oxygen atoms on there

so that would be called a resonance

hybrid structure

so we'll look at another example of of

resonance here using carbon dioxide

carbon dioxide is a great example

of resonance and so i will post

a video on on that as well

now when we're doing resonance

structures

sometimes we reach a

position where you can form a

two you know different resonance

structures uh between

different types of atoms and so you know

when we do

our carbon dioxide the double bonds or

your multiple bonds are going either

between carbon and oxygen or

carbon and oxygen well those are the

same in an earlier example

that we did for cobr2 we

saw that we put the double bond between

the carbon and the oxygen we did not

put any resonance forms between the

carbon and the bromine

now there's a reason that we did it that

way and the way that we figure that out

is using something called formal charge

i just said all those things so we'll

just skip those bullets for there

so what we really use formal charge for

is to determine which atom can best

handle a multiple bond okay it tells us

which structure is going to be preferred

by the molecule

uh when you have uh when you have

resonance not all resonance structures

uh are equivalent in energy and so

formal charge helps tell us which one

is going to be preferred now when we're

doing a formal charge calculation

on resonance structures uh it's it seems

like kind of a tedious process but the

math itself is extremely simple nothing

more than adding and subtracting

so when we're looking at formal charges

for resonance structures

we need to calculate a formal charge for

each atom in each possible structure so

if you have a large structure it's a lot

of calculations but again it's nothing

more than simple adding

and subtracting actually it's pretty

much all subtracting

so how do we do this so the first thing

that we do is we take a look at the

bonds that we have for

a single resonance structure for all of

the bonds that we have between

two atoms we split the electrons in the

bond

half of them get assigned to one of the

atoms that are in the bond and half of

them get assigned to the other atom

that are in the bond that's nice and

easy right

okay and then we look at the lone pairs

that we've assigned

to the atom okay and they belong to the

atom

that we put them on okay so if we look

at something like

water we saw that water had two lone

pairs on it

both of those lone pairs on the oxygen

belong only to the oxygen

and so we only count them for the oxygen

and so our final calculation for formal

charge is we take the number of valence

electrons that we expect

for a particular atom so something like

oxygen would be six

minus the number of electrons that we

have assigned it to the atom

in the molecule after we've split the

bonds and assigned the lone pairs

so we just take number of valence

electrons minus the number of electrons

we've assigned it

and that gives us the formal charge

okay and when you look at all of the

formal charges of

all of the different atoms in a

structure the preferred structure

is going to have all of the formal

charges as close as they can get

to zero now one thing you have to make

sure of that you calculated

to find that you've calculated your

formal charges correctly

is that the sum of all the formal

charges should add up

to the overall charge on the molecule if

they add up to something other than the

overall charge

on that molecule or or ion then you've

calculated your formal charge

incorrectly

okay but when you look at all of the

formal charges for all of the different

atoms

you want to see how close are all of

those charges to zero

and the preferred structure has the

charges closest to zero

okay now occasionally what you'll get is

you'll have more than one structure with

the same

set of formal charges but they might be

on

different atoms and if this is the case

then the structure

that has a negative formal charge on the

more

electronegative atom will be preferred

and i will show you an example of that

in a video okay and there's what i

mentioned earlier

okay so here's some examples we're going

to look at carbon dioxide

going we'll look at the ocn minus

ion and notice it says bonded in this

order so

in this case the oxygen's bonded to the

carbon carbons bonded to the nitrogen so

the

carbon here is actually the central atom

uh in this ion that's why i very

specifically said bonded in this order

if i ever uh give you a uh a molecule or

ion

where the first non-hydrogen atom is not

the central atom i will tell you

which one goes in the center or i will

give you the order in which they are

bonded

go back that real quick so look for

these examples done again as a separate

link

as a youtube video and then you can see

uh

how we use formal charge and how we

calculate it

now sometimes we will see molecules that

need

other than octets so we talked about

filling the octets of those outer atoms

filling the octets of the inner atoms

the octet is a good general rule but we

do have exceptions

for different elements anytime we have

beryllium which is actually fairly rare

but beryllium prefers to be surrounded

by four electrons rather than

eight it's kind of unusual but it's just

the way it goes

boron is also a strange strange element

it prefers to have

six it is uh capable of having eight

around it but it is actually more stable

when it just has

six electrons around it now these are

good just to kind of you know tuck in

the back of your mind for knowledge

this is not something i would test you

on i would not give you a

boron compound and then mark you wrong

because you didn't put

six instead you put eight okay so uh

again

tuck them in the back of your brain i am

not going to test you on

any types of exceptions like that now

other things to keep in mind is

any time you have a non-metal that is

phosphorous and

later in the periodic table it can form

additional bonds when we are doing

dot structures and these are something

that are called expanded octets when i

say additional bonds

that means that we can go over four

bonds to it so that means it can hold

more than eight

electrons it can also form uh

four bonds and still have some lone

pairs

which means it will have larger than an

octet

on it and here's a couple of examples we

have

arsenic penta fluoride you can see that

the arsenic

has has five uh bonds to it so it's got

ten electrons so we've got these guys

just right here so our one

two three four five so that's going to

be

10 electrons associated with that

arsenic

and for the sulfur we see the sulfur has

six

bonds to it each of those bonds count as

two electrons

to that sulfur and so that sulfur

actually has 12 electrons around it so

we call these

expanded octets

the other thing that we can also have

are odd numbered electron species

these are not terribly common but they

do happen

in certain chemical fields

especially if you choose to go into or

choose to study atmospheric chemistry

atmospheric chemistry occurs uh you know

you get stuff way up in the stratosphere

or higher

where they are exposed to some very high

energy sunlight some very high energy

ultraviolet radiation that we do not get

down here on the surface of the earth

and things start getting a little

strange

uh with some of that really energetic

light and you do get some odd numbered

species

up there some of those exist uh down

here as well but not nearly as common

well that finishes out the second

lecture uh i hope all this makes sense

make sure that you watch

the videos that have the examples in

there

and if you have any questions make sure

you contact me