CHEM 257 - Fall 2024 - Lecture 1 - Video 2

Summary

This chemistry lecture delves into molecular bonding, focusing on ionic and covalent bonds. Ionic bonding involves the transfer of electrons between a metal and non-metal, forming charged ions. Covalent bonding is further divided into polar and non-polar, determined by electronegativity differences and molecular geometry. The lecture also revisits Lewis structures, a fundamental tool for visualizing molecular connectivity, and introduces advanced concepts like multiple bonds and formal charge. The goal is to understand and predict molecular behavior through these bonding principles.

Takeaways

• 🔬 Ionic bonding involves a transfer of electrons between a metal (losing electrons, forming cations) and a non-metal (gaining electrons, forming anions), forming compounds like potassium chloride (KCl) and sodium chloride (NaCl).
• 🔍 Polar covalent bonds are based on the difference in electronegativity between atoms, leading to an unequal sharing of electrons and creating a dipole moment, as seen in molecules like HCl.
• 🌐 Non-polar covalent bonds occur when electrons are shared equally between atoms with a small difference in electronegativity, resulting in no net dipole moment, exemplified by methane (CH4).
• 📐 Molecular geometry and the interaction of lone pairs play a crucial role in determining the polarity of molecules, with symmetry often cancelling out dipole moments, as in the case of carbon dioxide (CO2).
• 📚 Electronegativity is a key concept in understanding the nature of chemical bonds, with the periodic table providing a reference for comparing the strength of the pull on a bound pair of electrons.
• 📝 Lewis structures are a fundamental way to represent molecular bonding, showing valence electrons and indicating possible molecular geometries, though they may not always be the best representation.
• 📋 Condensed formulas provide a shorthand notation for writing compounds, omitting bonding details but showing the total number of atoms and elements.
• 💠 Bond line or skeleton forms are a structural representation that shows how atoms are connected, omitting some atoms for clarity and focusing on key elements of the structure.
• 📉 The process of creating a Lewis Dot Structure involves determining basic connectivity, counting valence electrons, adding electrons to the most electronegative atoms first, and considering multiple bonds to satisfy the octet rule.
• ⚖️ Charged species require an adjustment in the total count of valence electrons to account for the charge, adding or subtracting electrons as necessary.
• 🔑 Formal charge and resonance structures are advanced concepts that will be used to determine the most stable and accurate representation of a molecule's structure in future lectures.

Q & A

What is the first type of bonding discussed in the script?

The first type of bonding discussed is ionic bonding, which involves a positively charged ion (cation) and a negatively charged ion (anion), formed by the transfer of electrons between elements.

What are some common examples of ionic compounds mentioned in the script?

Examples of ionic compounds mentioned include potassium chloride (KCl) and sodium chloride (NaCl), as well as potassium carbonate, which involves a metal and a polyatomic anion.

What is the significance of electronegativity in covalent bonding?

Electronegativity is significant in covalent bonding as it determines the difference in the pull of a bound pair of electrons, which can lead to polar or nonpolar covalent bonds.

What are the two main types of covalent bonds discussed in the script?

The two main types of covalent bonds discussed are polar covalent bonds, which have a difference in electronegativity, and nonpolar covalent bonds, where electrons are shared equally due to a small or no difference in electronegativity.

How does molecular geometry contribute to the polarity of a molecule?

Molecular geometry contributes to polarity by determining the spatial arrangement of atoms and the distribution of electron density, which can result in polar molecules if there is an uneven distribution of electron density.

What is the role of lone pairs in determining the polarity of a molecule?

Lone pairs are areas of high electron density that can affect the overall electron distribution in a molecule, potentially leading to a dipole moment and making the molecule polar.

Can you explain the concept of symmetry in relation to nonpolar molecules?

Symmetry in molecular structure means that the arrangement of atoms is such that any polarities are balanced and cancel each other out, resulting in a nonpolar molecule, as seen in the example of carbon dioxide (CO2).

What are the three structural representations mentioned in the script?

The three structural representations mentioned are Lewis structures, condensed formulas, and bond line or skeleton forms, each providing different levels of detail about molecular composition and bonding.

How is the basic connectivity determined in a Lewis structure?

The basic connectivity in a Lewis structure is determined by identifying which atoms are attached to each other, typically placing the least electronegative element in the center and connecting other atoms accordingly.

What is the octet rule, and how does it apply to the Lewis structure of the molecule discussed in the script?

The octet rule states that atoms tend to form bonds in such a way that they have eight electrons in their valence shell, achieving a stable electron configuration. In the script, the carbon atom in the molecule COCl2 initially does not have a full octet, leading to the formation of a double bond with oxygen to satisfy the octet rule.

How do you determine the correct number of valence electrons for a molecule or ion?

The correct number of valence electrons is determined by counting the valence electrons from each atom and adding them together. For charged species, you must also account for the charge by adding or subtracting electrons accordingly.

Outlines

00:00

🔬 Ionic and Covalent Bonding Basics

This paragraph introduces the fundamental concepts of ionic and covalent bonding in chemistry. Ionic bonding is explained as the process where a metal (cation) and a non-metal (anion) form a compound by exchanging electrons. Examples such as potassium chloride and sodium chloride are given. The importance of electron transfer in forming ionic compounds is emphasized. Covalent bonding is then divided into two types: polar and non-polar. Polar covalent bonds are characterized by a difference in electronegativity between atoms, leading to an unequal sharing of electrons. Molecular geometries and the interaction of lone pairs are highlighted as key factors in determining polarity. Non-polar covalent bonds involve equal sharing of electrons with minimal electronegativity difference, resulting in no net dipole moment. The paragraph sets the stage for a deeper exploration of these concepts in subsequent lectures.

05:01

📚 Understanding Polar and Nonpolar Covalent Bonds

The second paragraph delves deeper into the nature of polar and nonpolar covalent bonds. It explains that the polarity of a bond is determined by the difference in electronegativity between the atoms involved. A visual representation using arrows to indicate electron density towards the more electronegative atom is introduced. The concept of molecular geometry and its impact on the overall polarity of a molecule is discussed, using carbon dioxide as an example of symmetrical cancellation of dipole moments. The paragraph also contrasts polar and nonpolar bonds, highlighting that nonpolar bonds occur when the electronegativity difference is minimal and the electrons are shared equally. Methane is presented as a classic example of a nonpolar molecule due to the small electronegativity difference between carbon and hydrogen. The importance of symmetry in determining molecular polarity is underscored, and the paragraph concludes with an introduction to different structural representations of molecules, such as Lewis structures, condensed formulas, and bond line or skeleton forms.

10:02

📝 The Process of Creating Lewis Dot Structures

This paragraph outlines the method for creating Lewis dot structures, starting with the basic connectivity of atoms. The least electronegative element is placed in the center, and other atoms are attached accordingly. The process involves counting valence electrons from each atom and adjusting for the molecule's charge. The importance of electronegativity in determining which atoms receive electrons first is emphasized. The paragraph explains the rule of placing electrons on the most electronegative atoms and forming multiple bonds when necessary to satisfy the octet rule for atoms like carbon, nitrogen, oxygen, and fluorine. The example of Cl2CO is used to illustrate the step-by-step process of creating a Lewis structure, including the formation of a double bond to satisfy the carbon's need for a full octet. The paragraph concludes with a brief mention of resonance structures and formal charge, topics that will be revisited in future lectures.

15:03

🔍 Advanced Lewis Structure Construction

The final paragraph provides an in-depth look at the Lewis structure construction process, focusing on the distribution of valence electrons and the formation of multiple bonds. It explains the rule of prioritizing the most electronegative atoms when placing electrons and the subsequent need to form multiple bonds to complete the octet for certain atoms. The example of Cl2CO continues, demonstrating how to address the carbon's incomplete octet by forming a double bond with oxygen, which involves the donation of a lone pair from the oxygen atom. The paragraph also touches on the concept of resonance structures and the use of formal charge in determining the most stable Lewis structure. It concludes by summarizing the importance of understanding different types of bonds and the process of constructing Lewis structures, setting the foundation for further studies in chemistry.

Keywords

💡Ionic Bonding

Ionic bonding is a type of chemical bond formed through the electrostatic attraction between oppositely charged ions, typically involving a metal cation and a non-metal anion. In the video, ionic bonding is discussed as the first type of bonding, with examples such as potassium chloride and sodium chloride. The process involves the transfer of electrons from one atom to another, resulting in the formation of ions that are attracted to each other due to their opposite charges.

💡Covalent Bonding

Covalent bonding is characterized by the sharing of electrons between atoms, usually non-metals, to form molecules. The video distinguishes between two types of covalent bonds: polar and non-polar. Polar covalent bonds occur when there is a difference in electronegativity between the atoms, leading to an unequal sharing of electrons, while non-polar covalent bonds involve an equal sharing of electrons due to similar electronegativity values.

💡Electronegativity

Electronegativity is a measure of the tendency of an atom to attract a pair of electrons in a chemical bond. It plays a crucial role in determining the polarity of covalent bonds. In the script, electronegativity differences are used to explain the formation of polar covalent bonds, where one atom has a stronger pull on the shared electrons than the other.

💡Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. It is important in understanding the polarity of molecules because the shape can influence the distribution of electron density. The video mentions molecular geometry in the context of how it, along with electronegativity and lone pairs, contributes to the polarity of covalent compounds.

💡Lone Pairs

Lone pairs are pairs of electrons that are not involved in bonding and are located on an atom. They can influence the shape and polarity of a molecule. In the script, lone pairs are discussed as a factor that can affect the polarity of a molecule, especially when they interact with the molecular geometry and the distribution of electron density.

💡Polar Molecules

Polar molecules are molecules that have an uneven distribution of electron density, resulting in a molecule with a positive and a negative end, known as a dipole moment. The video explains that polar covalent bonds, molecular geometry, and lone pairs can all contribute to the polarity of a molecule, giving examples such as carbon dioxide, which despite having polar bonds, is non-polar due to its symmetrical shape.

💡Non-polar Molecules

Non-polar molecules are those in which the electron density is distributed evenly, resulting in no net dipole moment. The video provides the example of methane, where the electronegativity difference between carbon and hydrogen is small, leading to an equal sharing of electrons and a lack of a net dipole moment.

💡Lewis Structure

A Lewis structure is a graphical representation of a molecule that shows the arrangement of atoms using lines to represent covalent bonds and dots to represent valence electrons. The video script revisits the process of creating Lewis structures, starting with the basic connectivity and then counting valence electrons to determine the bonding and lone pairs in molecules like Cl2CO.

💡Condensed Formula

A condensed formula is a shorthand representation of a chemical compound that omits the bonding lines and shows only the types and numbers of atoms present. It is useful for quickly writing out complex molecules. The video mentions condensed formulas as a structural representation that will be discussed further in a subsequent lecture.

💡Bond Line Form

Bond line form, also known as skeleton form, is a type of structural representation that shows how atoms are connected in a molecule, typically omitting hydrogen atoms for clarity. The video introduces bond line form as a new concept for students, which will be practiced and explained in more detail in a later lecture.

💡Octet Rule

The octet rule states that atoms tend to form bonds in such a way that they each have eight electrons in their valence shell, giving them the same electronic configuration as a noble gas. In the script, the octet rule is mentioned in the context of forming multiple bonds to satisfy the electron requirements of atoms that are short of eight valence electrons, such as the carbon in the Cl2CO example.

Highlights

Introduction to the two types of chemical bonding: ionic and covalent.

Explanation of ionic bonding involving cations and anions formed by electron transfer.

Examples of ionic compounds like potassium chloride and sodium chloride.

Discussion on metal can species and polyatomic anions in ionic compounds.

Polar covalent bonds based on electronegativity differences and molecular geometries.

Importance of electronegativity in determining the polarity of covalent bonds.

The concept of nonpolar covalent bonds where electrons are shared equally.

Illustration of polar and nonpolar bonds using molecular geometry and symmetry.

Process for determining if a molecule is polar by analyzing electronegativity differences, molecular geometry, and lone pairs.

Introduction to Lewis structures and their role in depicting molecular bonding.

Guidelines for creating Lewis structures, including basic connectivity and electron counting.

Explanation of how to account for charged species in Lewis structures.

Step-by-step process for drawing Lewis structures for a given molecule.

The significance of the octet rule and multiple bonding in Lewis structures.

How to determine which atom will donate lone pairs to form multiple bonds.

Different structural representations: Lewis structure, condensed formula, and bond line or skeleton form.

Importance of symmetry in determining the overall polarity of a molecule.

Review of the process for creating Lewis Dot Structures for a specific molecule, Cl2CO.

Final remarks on the importance of understanding different types of bonds and Lewis structures in chemistry.

Transcripts

hi students for the second video for

lecture one we're now going to continue

our general chemistry review and now

talk about the two types of bonding so

remember that we talked about our

molecular bonding in chapter 3 in your

general chemistry one course so what is

the first type of bonding that we want

to talk about well that is going to be

our ionic bonding remember that involves

a positively charged ion which we would

commonly call A

cation and a negatively charged ion so

the important part that you want to have

for the ionic compounds is that you must

form the ion this is done by gaining or

losing electrons from our elements

various examples that you would see back

in your general chemistry were such

things as potassium I chloride you would

also see such things as sodium chloride

and then you got to species that you

really didn't like and that is your

metal cans so in this case pottassium

but then your polyatomic anions an

example of that being such things as

carbonate so this being potassium

carbonate so when dealing with ionic

compounds it was normally a metal and a

non-metal one losing an electron the

other gaining it and forming a compound

based off of the charge difference this

is how ionic bond works for Cove valent

Bond we actually are going to have two

types of bonds that we need to

discuss the first of which is the polar

calent polar calent bonds are going to

be based off the difference of one of

our periodic trends that you really

needed to focus in back on your general

chemistry this is going to be the

difference in electr

negativity remember that electro

negativity is a relationship Trend in

which we are comparing the strength of

the pull of a bound pair of electrons in

a CO valent polar or non-polar Bond so

this is going to specifically talk about

the tug of war on a pair of electrons in

a bond so the question becomes how or

why does this all work well the first

thing that we need to really think about

in here is that our difference in electr

negativity

then mixed with our ideas of molecular

geometries as well as the interaction

with lone pairs

is what's going to give us our possible

polar species three very very very

important ideas that all lead to our

polar species so it's electr negativity

difference our molecular geometries then

inside of our molecular geometries the

interactions with our lone pairs finally

giving us the possibility of polar Cove

valent compounds we will talk about how

to study that in our next lecture

finally the other type of Cove valent

bond is the non-polar so what is a

nonpolar Bond well if a polar bond is

talking about the difference in electro

negativity for a non-polar Bond we are

going to share our electrons

equally this means that the difference

in electr negativity is very small and

those atoms in the bond are going to

have an equal right to those electrons

so we don't have a electron moment going

towards either one of the atoms so what

does this actually look like for us well

as we're about to see with our polar

species if I have something like

HCL we are always going to have a really

nice way of indicating which species is

going to be more polar we are going to

draw an arrow in which the arrow is

pointing towards the more electr

negative atom and a little plus sign is

going to be here above the more

electropositive atom that way you can

see that the electrons for this HCL Bond

are going to go more towards the

chlorine so in a good example of this

and talking about our molecular

geometries is if I have oxygen bound to

carbon bound to another oxygen I'm going

to form carbon

dioxide but in this case one of the

oxygens is going to pull electron

density towards that oxygen while the

other oxygen is going to pull electron

density towards itself notice that these

are going to cancel out this is

typically what we call symmetry we are

going to see a lot of these examples

when we get to our second lecture so

what does it mean for a nonpolar

compound to share electrons equally well

if I do that probably the best example

is going to be methane this is carbon

surrounded by four

hydrogens the electr negativity

difference between carbon and hydrogen

is not that large so therefore it's

going to be considered a very equal

sharing and so all of the possible

dipole moments if there are any would a

cancel out but there aren't going to be

any due to the lack of a noticeable

electr negativity difference it's really

important that we go through a very nice

seamless process in order to determine

whether something is polar or not the

best way of doing that is start with the

electr negativity difference of all of

our bonds

so this means that we are going to need

to use our periodic trend and probably

determine the calculated value in order

to see whether if an individual bond is

polar or not and then once we have done

that then we will go to the molecular

geometry I'm going to abbreviate this

just as mg molecular geometries are

going to show us that certain molecular

geometries are going to lend us to the

possibility of certain polar molecules

we will talk about those in lecture two

finally once we get done with the

molecular geometries we must go and see

if there are lone

pairs on our Central atoms lone pairs

are going to be areas of high electron

density and might lead us to having a

possible dipole negative portion of our

molecule and then finally one area that

we must go and look at is our symmetry

this was seen in the carbon dioxide

example listed up above carbon and

oxygen individually are going to be

considered polar bonds but because

they're pulling opposite and equal you

are seeing that this is going to be a

nonpolar molecule that contains bonds so

for our blue example above that's not a

polar molecule because it has dipole

moments pulling in opposite and equal

directions so now the question becomes

how can I represent all of the various

different structures that we're going to

see in this course well there are going

to be three that we will run into all

the time you know the first one this is

going to be the Lewis structure the

Lewis structure is the type of molecular

representation that we used in chem 119

the Lewis structure is going to take the

veence electrons will show very

rudimentary bonding and give us a little

bit of a possible indication when it

comes to the different molecular

geometries but this may not be the best

representation of our structure another

way of doing this is what we call a

condensed formula a condensed formula is

going to no longer show any bonding but

is going to be a very nice shorthanded

information about what atoms I have

bound to one another and going to just

show a total number of atoms and

elements in our compound this is

actually very nice when we are writing

out compounds and we were trying to put

them in paragraph form finally probably

the most important of

our structural representations is going

to be a new one for you and one that we

will practice through chapter one and

that goes by a couple of different names

probably the most common name is known

as Bond line

form but many textbooks will also call

this a skeleton form skeleton form

meaning that it's going to show you how

things are

connected and it is going to show some

very important certain atoms while going

and omitting some of the atoms for a

little bit of clarity sake so as you can

see as we go along you should be very

good with the Lew structure but I will

show you how to do that here on the next

slide while we will talk about the

condensed formula and the bondline

skeleton form when we get to lecture

two so now let's go back and review a

simple process for creating a Lewis Dot

Structure and we're going to go ahead

and do that for

cl2

Co so what we need to do in here is

first start with the basic connectivity

we need to go and figure out what atoms

are attached to each other the best way

of explaining this is that the least El

electr negativity element should be

located in the middle now what I will do

is attach an oxygen to it and I will

attach one chlorine to either side this

is the basic skeleton structure what is

so important about this is that it is

going to give us our basic

connectivity so now the question becomes

what do we do next now we need to count

the veence electrons coming from the

periodic table so when we are counting

electrons it is important to remember

that we are only using the veence

electrons because those are the ones

that are going to help us with bonding

so for our chlorines you will notice

that I have two of them they are in the

hallogen column which is going to give

us seven electrons so we will multiply

those by two chlorines and we will get a

total of 14 electrons from chlorine

alone when it comes to my carbon I am

going to get four

electrons and finally with my oxygen I

am going to get six veence

electrons and now all I'm going to do is

add this up if I take the sum of all of

my four atoms you will see that I get

24 veence

electrons this is really nice because

this is a neutrally charged species but

we must go back and review a rule that

we discussed back in our general

chemistry when we were talking

about our individual species so now with

that being the case we have to go and

look at a negatively charged species

such as NO3 1

minus what we need to do is count count

the total number of electrons I get five

from the nitrogen I get six from each of

the oxyg so if I do 6 * 3 is 18 + 5 is

23 however the correct number of

electrons for this compound is actually

24 electrons the reason why is that we

must be concerned with this negative

charge located here in is an additional

electron that we must take account for

so when we have our charged

species we must go and add or subtract

electrons due to the charges so if you

see a plus charge on your ions something

like ammonia this means that it is going

to be more positive than it expects so

you are going to

lose at least one

electron so now what will happen when I

have a negative charge well a negative

charge just as we have seen here with

our nitrate we are going to have to

add an electron so this is really

important that you go and get the

correct number of electrons so that way

when you are doing your leis dot

structures you are coming up with the

right number to help you build the

structures that you

need so now we need to go and draw in

our bonds the best rule for when adding

electrons to our skeleton structure is

that we are going to add electrons to

the most electronegative atoms first so

in this case I am going to go and redraw

my skeletal

structure and now my oxygen and

chlorines are going to get the remaining

electrons the question becomes how many

remaining electrons do I have well I

start with 24 electrons but based off of

the basic connectivity I have to go and

subtract six

electrons and now it is my job to place

18 electrons around my atoms so let us

go and do that to chlorine first the

left chlorine now has six the chlorine

here now gets six more as lone Pairs and

finally six more when it comes to the

oxygen we now have a new problem and

that is my carbon does not have a full

octet so this will now lead us into our

fourth step the fourth step is now

talking about the possibility of

multiple bonds this is going to happen

if we have

atoms with

fewer than eight

electrons this is commonly known as the

octet rule if we are dealing with

certain species especially the ones that

we are going to use in organic chemistry

you will see a lot that must obey the

octet rule carbon nitrogen oxygen

and Florine so our carbon is currently

two electrons short so what we are going

to do is go and form multiple Bonds in

order to help out that compound's

current structure and electronic

Behavior so the first thing that we can

think about is going and forming a

double bond remember that a double bond

is going to be four electrons in a

shared Bond that's not the only only one

that we can do we are also able to go to

the triple

bond that is going to be six electrons

shared

coal so now if we go back to our third

step we need to figure out which one our

chlorine or our oxygen is going to go

and donate one of its loan pairs to help

the carbon form a double bond the best

way of explaining this is the species

that have the lone pairs you are going

to donate from the one that is less

electronegative in this example it's

going to be the oxygen so I'm going to

take one of the pairs of Lone pairs of

electrons on my oxygen and I'm now going

to go in and form the multiple Bond what

my final structure is going to look like

is I'm going to have a double bond to my

oxygen and then I will have single Bond

bonds to my chlorines in which the

chlorines then have three pairs of Lone

pairs of electrons you will see in later

lectures just how we are able to

determine which structures are better we

typically call those resonance

structures and we are going to use such

rules as formal charge that we learn

back in our general chemistry

course in this video you just saw how

important it was to discuss the

different types of bonds not only ionic

bonding which uses charge different but

then the two types of calent bonds polar

and non-polar finally we went back and

Revisited an old friend from General

chemistry and we started our review of

the lewis. structure which we will use

to then build further structures in

future

classes have yourself a great day