6.2 Covalent Bonding and Molecular Compounds

Summary

This video discusses covalent bonding and molecular compounds, explaining how atoms rarely exist independently and instead form molecules through covalent bonds. The video highlights key concepts such as molecular formulas, bond energy, the octet rule, Lewis structures, and multiple bonds. Examples like hydrogen, fluorine, and nitrogen bonding are provided, illustrating single, double, and triple bonds. The video also explains how resonance structures, such as ozone, exist as hybrids of multiple possible structures. Overall, it focuses on the principles that govern the formation and stability of molecular compounds.

Takeaways

• 😀 Covalent bonds involve atoms sharing electrons to form molecules, which are electrically neutral.
• 🔬 Molecular compounds consist of atoms bonded together, and they are represented by molecular formulas.
• 💧 Diatomic molecules are made of two atoms, like hydrogen (H2) and fluorine (F2).
• ⚛️ Atoms bond to reduce their potential energy, forming stable compounds at a specific bond length.
• 🌱 Bond energy refers to the energy required to break chemical bonds and is measured in kilojoules per mole.
• 🔋 The octet rule states that atoms tend to form compounds to achieve eight electrons in their outermost shell.
• 🧪 Single bonds involve one pair of electrons, while double and triple bonds involve multiple electron pairs, making them stronger and shorter.
• 📐 Lewis structures represent the arrangement of electrons in a molecule, showing bonded and unbonded electron pairs.
• 🧬 Resonance structures exist when a molecule can be represented in multiple ways, as with ozone (O3).
• 💥 Multiple bonds (double or triple) are shorter and stronger than single bonds due to closer nuclei in shared orbitals.

Q & A

What is a covalent bond?

A covalent bond is a type of chemical bond where two or more atoms share electrons, creating a molecule that is electrically neutral.

What is a molecular compound?

A molecular compound consists of molecules formed by atoms that are bonded together through covalent bonds. It has a specific molecular formula that indicates the number and type of atoms in the compound.

What is a diatomic molecule, and can you give an example?

A diatomic molecule is composed of only two atoms, which can be of the same or different elements. For example, hydrogen gas (H₂) and fluorine gas (F₂) are diatomic molecules.

Why do atoms form molecules instead of existing as standalone atoms?

Atoms tend to form molecules because bonded atoms have lower potential energy compared to standalone atoms. Lower potential energy makes the molecules more stable, which is a preferred state in nature.

What is bond energy, and why is it important?

Bond energy is the energy required to break a chemical bond and separate atoms in a molecule. It is important because it represents the strength of the bond; the same amount of energy is released when the bond is formed as is required to break it.

What is the octet rule?

The octet rule states that atoms tend to form bonds in such a way that each atom has eight electrons in its outer shell, achieving the same electron configuration as a noble gas, which is a stable, low-energy state.

How does the bonding in hydrogen differ from the bonding in larger atoms like fluorine?

Hydrogen forms a covalent bond by sharing its single electron with another atom, aiming to achieve a stable 1s² configuration (like helium). Fluorine, however, shares electrons to complete its octet, filling its 2s and 2p orbitals.

What is the significance of bond length, and what is the bond length for hydrogen?

Bond length is the distance between two bonded atoms at their lowest potential energy. For hydrogen, the bond length is about 75 picometers, where the attractive forces between electrons and nuclei balance the repulsive forces.

What are multiple bonds, and how do they differ from single bonds?

Multiple bonds are covalent bonds where two or more pairs of electrons are shared between atoms. A double bond shares two pairs, and a triple bond shares three pairs. Multiple bonds are shorter and stronger than single bonds, which share only one pair of electrons.

What is resonance, and why does it occur in molecules like ozone?

Resonance occurs when a molecule cannot be represented by a single Lewis structure because the real structure is a hybrid of two or more forms. In ozone (O₃), the bonds between oxygen atoms exist as a blend of single and double bonds, leading to resonance.

Outlines

00:00

🧬 Covalent Bonding and Molecular Compounds

In this video, the concept of covalent bonding and molecular compounds is introduced. Atoms rarely exist independently and usually form molecules, which are neutral groups of atoms held together by covalent bonds. This contrasts with ionic bonds, which do not form molecules. Covalent bonds involve the sharing of electrons between atoms, leading to the creation of molecular compounds. These compounds are represented by molecular formulas, showing the types and numbers of atoms. For instance, water (H2O) consists of one oxygen atom and two hydrogen atoms. Additionally, diatomic molecules, like two hydrogen atoms bonding together, are introduced.

05:02

⚛️ Potential Energy and Bond Formation

This section discusses how atoms, such as hydrogen, have higher potential energy when separated and lower potential energy when bonded. As atoms approach each other, the attraction between their nuclei and electrons grows, surpassing the repulsive forces. Eventually, they reach a distance called the bond length, where potential energy is minimized, and the atoms form a stable bond. For hydrogen, this bond length is about 75 picometers. The energy released during bond formation is called bond energy, and breaking the bond requires the same amount of energy. For hydrogen, the bond energy is 436 kilojoules per mole.

10:03

🔗 Octet Rule and Covalent Bonding

The octet rule states that atoms tend to form compounds to achieve eight electrons in their outer shell, similar to noble gases. This section explains how atoms like fluorine bond to achieve a stable octet. Fluorine atoms share electrons to achieve eight valence electrons. Similarly, hydrogen and chlorine can share electrons to complete their outer shells, mimicking noble gas configurations. The concept of the octet rule is crucial in understanding why atoms form chemical bonds, and electron sharing is a key aspect of covalent bonding.

15:05

📝 Lewis Structures and Structural Formulas

Lewis structures are introduced as a simpler way to represent the arrangement of electrons in bonded atoms. In these diagrams, the atomic symbol represents the nucleus and inner electrons, while dots show valence electrons. Bonds are represented by lines, which denote shared pairs of electrons. The section explains how these structures can illustrate single, double, or triple bonds between atoms, with each type of bond representing different numbers of shared electron pairs. It also mentions that structural formulas, which use lines to represent bonds, are more practical for larger molecules.

🌍 Resonance Structures and Chemical Bonding

This section addresses the concept of resonance, where some molecules, like ozone (O3), cannot be represented by a single Lewis structure. Instead, they are a hybrid of multiple structures, which are shown with resonance arrows between them. This means that the actual bonding in these molecules is a combination of the possible structures, and the molecule does not exist as either structure independently. Resonance structures help accurately depict the behavior of certain molecules that don't conform to a single electron arrangement.

Keywords

💡Covalent Bond

A covalent bond is a type of chemical bond where two atoms share one or more pairs of electrons to achieve stability. In the video, covalent bonds are central to the formation of molecules and molecular compounds, allowing atoms to reach a lower potential energy state. For example, two hydrogen atoms share electrons to form a stable H2 molecule.

💡Molecule

A molecule is a group of two or more atoms held together by covalent bonds. The video emphasizes that molecules are electrically neutral and can be composed of the same or different atoms. An example provided is the water molecule (H2O), where two hydrogen atoms and one oxygen atom form a molecular compound.

💡Molecular Formula

A molecular formula provides the types and numbers of atoms in a molecule. For example, the molecular formula for water is H2O, meaning the molecule consists of two hydrogen atoms and one oxygen atom. The video explains that this formula helps represent molecular compounds accurately.

💡Diatomic Molecule

A diatomic molecule consists of two atoms, either of the same or different elements. The video introduces diatomic molecules by describing elements like hydrogen (H2) and fluorine (F2) that naturally form diatomic structures. The term 'diatomic' reflects the combination of 'di-' meaning two and 'atomic' referring to atoms.

💡Potential Energy

Potential energy refers to the energy stored in an object due to its position or arrangement. In the context of covalent bonding, atoms start with high potential energy when separated, which decreases as they approach each other to form a bond. The video explains how atoms achieve lower potential energy through bonding, reaching a stable configuration.

💡Bond Energy

Bond energy is the amount of energy required to break a bond between two atoms. It is also the energy released when a bond forms. In the video, bond energy is measured in kilojoules per mole and helps explain the stability of diatomic molecules like hydrogen (H2), which require 436 kJ/mol to break their bonds.

💡Bond Length

Bond length is the distance between the nuclei of two bonded atoms where the potential energy is at its minimum. The video describes bond length in the example of hydrogen molecules, where the bond length is about 75 picometers. At this distance, the attractive and repulsive forces between atoms are balanced.

💡Octet Rule

The octet rule states that atoms tend to form compounds in such a way that each atom achieves eight electrons in its outermost shell, achieving stability similar to noble gases. The video illustrates this with examples like fluorine and chlorine, where atoms share or transfer electrons to fulfill the octet rule.

💡Lewis Structure

A Lewis structure is a diagram that represents the bonding between atoms in a molecule, showing how valence electrons are shared or paired. In the video, Lewis structures are used to visualize covalent bonds, such as in the diatomic fluorine molecule, where shared electrons are shown as dots or lines between atoms.

💡Resonance

Resonance refers to a situation where a molecule can be represented by two or more valid Lewis structures, with the true structure being a hybrid of these forms. The video explains resonance using ozone (O3) as an example, where the bonding arrangement cannot be captured by a single Lewis structure but instead fluctuates between two forms.

Highlights

Introduction to chapter six, section two on covalent bonding and molecular compounds.

Atoms rarely exist in nature as standalone entities; they usually form molecules.

Molecules are electrically neutral and held together by covalent bonds.

Covalent bonds involve the sharing of electrons between atoms.

Molecular compounds have molecular formulas that represent the number and type of atoms.

Diatomic molecules consist of only two atoms bonded together, like hydrogen (H2).

Atoms joined together have lower potential energy compared to when they are isolated.

Bond length is the specific distance where attraction and repulsion between atoms balance, and for hydrogen, this length is 75 picometers.

Bond energy is the energy required to break a bond, measured in kilojoules per mole.

The octet rule states that atoms tend to form compounds to achieve eight electrons in their valence shells.

Fluorine forms diatomic molecules (F2) by sharing electrons to satisfy the octet rule.

Lewis structures represent covalent bonds and unshared electron pairs.

Structural formulas simplify larger molecules by focusing on bonds without showing all electron pairs.

Multiple bonds (double or triple) are shorter and stronger than single bonds.

Resonance structures, such as ozone (O3), represent molecules that exist as hybrids of different bond arrangements.

Transcripts

so in this video we'll be covering

chapter six section two which is

covalent bonding and molecular compounds

now as i mentioned previously atoms

rarely exist in nature as standalone

objects so what they'll usually do is

form what are called

molecules

now these molecules are electrically

neutral once they've been bonded

and they're groups of two or more atoms

and they're all held together by

covalent bonds

which is very important because

ionic bonds don't form water

technically called molecules so

basically what happens is that

you'll have

a few atoms

and they're sort of sharing electrons

in a covalent bond as we mentioned

before

and when these atoms

bond together they form what are called

molecular

compounds and these compounds

don't only have an illustration as i

have here they also have what is known

as a molecular formula

which

gives the number of atoms and what type

of atom are in each type of molecule so

for example let's say this was water

what you do is you would

take this oxygen here

and put it in the formula and then you

would also take the

two hydrogens

which are right here

i'm sure most of you know this formula

because it's a

very common phrase h2o

but what you do is you take the number

of the atom in this case two and then

the default if it's one you just leave

it blank so you'd get hydrogen

to oxygen

now if you just had

two hydrogen atoms

off to the side

bonded together

they would form what is known as a

diatomic

compound

and this is because

di means two and atomic means obviously

atoms

so a molecule that has only two atoms is

known as diatomic and we'll find a list

of

elements

that stand alone and create diatomic

atoms in nature

so as i explained in the last video

atoms that stand alone

tend to have a higher potential energy

than when they are joined together

with other atoms and now i'm going to

explain why and the way we're going to

do this is by visually

two

hydrogen atoms like this now

if they were very far apart without

influencing each other there'd be a

great potential energy for them to come

together

and form a compound now as they get

closer

what you'll find is that

the attraction between

the nuclei in each one to the other

atoms electron

is much higher than the repulsion

between these two

so what they'll do is they'll move

closer and closer together

building up momentum

and

changing their potential energy to

kinetic energy

as they move closer

again nature wants to move to a lower

potential energy state so what happens

is that these will keep moving together

until eventually

they reach

a magical distance where they are

in the lowest potential energy state

possible

so we'll call this e low

and this is the point where the

attraction

between the nuclei of each one

and the other one's electron

balances out the repulsion between

the nuclei and the

electrons and this is the lowest

potential energy state because if you

were to force them

even closer together

where they were almost on top of one

another

what you would find is that

this repulsion between the two nuclei

would be

so great that it would get rid of

potential energy

and would store it more in electrical

energy trying to force the

molecule apart so at this point where

the two hydrogen atoms

are bonded together

uh

and their potential energy

is at a minimum

the electrons in each atom

can orbit freely in either orbital

because they're sort of overlapped at a

minimum

energy state so the electrons have the

ability to

go from one atom's influence to the

other without doing any work

now this low energy state here

occurs consistently at a specific

distance

called the bond length

and in the hydrogen

the bond length is about 75

picometers

which means that once you get to this

point

the atoms will still vibrate a little

going from attraction to repulsion and

vice versa however uh

once the atoms are this far apart

they get to this covalent state where

the electrons can flow freely from one

to another and for those of you who know

about the law of conservation of energy

you may be wondering where all this

potential energy

that uh separated the atoms before has

gone and

it has gone into a form of energy known

as bond energy

now bond energy

is useful to know

because it is the same

going in as it is coming out meaning

that

the same amount of potential energy that

these atoms had when they were

way far apart uh

is the same amount that is released

when you break these atoms up

and this bundt energy which is again the

energy required to break

a chemical bond and make these atoms

neutral each with one

electron is measured in kilojoules

per mole

meaning they measure the amount of

energy it takes to completely break the

bonds in one mole of substance now the

bond energy for diatomic hydrogen like

the example we have over here

is 436

kilojoules

per mole meaning it takes

436 kilojoules of energy to break up

6.02 times 10 to the 23rd

uh bonds within these various molecules

and i know this is kind of the simplest

example with two of the most

two atoms of the most basic element

there is however these principles all

apply to the rest of covalent bonds

and just to further reiterate the

stability of this hydrogen to hydrogen

bonding i'll draw a little diagram of

what's happening in each

hydrogen's 1s orbital which is the only

orbital they

possess of course so they start off each

with

one electron in the orbital

but then once they bond you end up with

two hydrogens with a sort of shared

orbital

that has

one

electron of each spin and this gives it

the 1s2

configuration

which is of course what the noble gas

helium has and again

if you'll remember the noble gases have

the lowest potential energy within their

orbitals

which is why

hydrogen

bonds together like this in order to get

this 1s2 configuration which is the low

potential energy of noble gases

now the noble gases have this low

potential energy because their outer

orbitals their valence electrons

have completely filled

their s

and p orbitals or in the case of helium

uh just the s orbital

and these full s and p orbitals each of

which can hold two and six electrons

respectively

uh allow

the noble gases to have eight valence

electrons

now unfortunately for the rest the

periodic table they do not come with

eight valence electrons however they

still want to get to the state because

it is the lowest potential energy so

what they will do

is either share

give

or take

someone else's electrons

in order to get to this

eight electron configuration and this is

what is known as the

octet rule

now the octet rule says

that chemical compounds will tend to

form

so that

each atom

will have an outer

valence

of eight electrons

in its outermost shell so just to give

you an example we'll look at how two

atoms of independent fluorine bond to

form

a diatomic f2 now fluorine is a halogen

which means it has

seven valence electrons

given

in its 2s

orbital

and 2p orbital

you'll notice the two and the five add

up to seven

however if you

sort of separate

this last electron

and examine it next to

another fluorine atom

what you'll find

is that

if these two atoms sort of exchange

these electrons

so that at some points this fluorine

atom over here

will take over this electron

it will contain

eight electrons

in its

s p shells

and then this one at some points will

also contain eight giving it a stable

octet at some points

when they're close enough

to share

these outer electrons and the same thing

goes for the chemical

hcl which is one hydrogen and one

chlorine now if you look at the

arrangement of chlorine

its valence is in the third energy level

so it's 3s

is full

and it's 3p

once again because it's halogen has

five electrons

and we'll leave this last one off to the

side but know that it is in the 3p

orbital

and then if you look at hydrogen

which just has the 1s

orbital with the one

electron

now if you look if these two

share this electron

you'll see that chlorine will then have

eight total electrons in this shared

orbital

and hydrogen will have two

this gives chlorine the arrangement of

argon

and hydrogen the arrangement of helium

both of which are noble gases

and it gives chlorine this octet rule

hydrogen doesn't follow it because it

can't have a p orbital

but it does still form a stable noble

gas configuration

so now i'm going to be demonstrating a

much easier way of representing uh

an element's electrons without having to

write out the full 1s 2 2 s2 etc

electron configuration notation instead

what you can do is you can take an

element

let's say element x and you can just dot

how many

valence electrons

are around it up to eight

so if we go across the second period and

do this you'll find lithium with one dot

in its valence beryllium two

boron

three

carbon four

nitrogen

five

oxygen six

fluorine seven

and finally neon has the full octet

with eight

and this notation can be very useful for

illustrating bonds

for example if we take the fluorine

fluorine bond that we did earlier and we

draw out

the electron dot notation

there's seven on that fluorine and seven

over here on this fluorine

you can put them together and see that

these two right here are a shared pair

giving each

eight the full octet

independently

now this shared pair can also be

represented

by a line so we could alternately draw

this as f

with

its seven electrons

the other fluorine with its seven

electrons

covalently bonded represented by this

line now in this instance there is the

one bonded pair

of electrons in the middle represented

by the line and the rest of these are

what are known as unbonded pairs of

electrons meaning they aren't involved

in the bond between the two atoms and

drawings like this are what are known as

lewis structures now lewis structures

as i mentioned earlier are things

drawings where the atomic symbol in this

case f represents the nucleus

and all the inner shell electrons

the non-valence electrons that is

and then the dashes represent a covalent

bonds

but it's not uncommon to leave off the

unshared electrons that aren't involved

in the bond so for example

for the third example of how this

fluorine bond could be represented you

could just do

f dash f and that would represent the

diatomic fluorine and all chemists would

know that there's six unshared electrons

on each fluorine

now this is not a lewis structure here

the lewis structure shows all the

electrons this is what is known as a

structural formula

and this becomes much more practical

when you get into

much larger

molecules

and you have

you know 10 or 20 different atoms bonded

you don't want to be putting all these

dots along you just want to show how the

atoms are bonded to one another within

the case of fluorine where only one pair

of electrons

is bonded as well as other things like

uh

hydrogen or

hydrogen chloride

these are what are known as a single

bond

molecules because only one pair

of electrons is being shared atoms

aren't always simply bonded

one to another by

single bonds however

some atoms for example carbon

can form what are called

double bonds where two pairs of

electrons are shared which can also be

written as two different dashes

or nitrogen for example can even form

triple bonds

and this is because if you look at the

electron dot notation for nitrogen which

has five electrons in its valence

what you'll find

is that if you share

this top pair

the middle pair

and the bottom pair

is that each one in this shared orbital

then has a grand total of six in the

shared plus the two it already has to

form the octet rule

only by following triple bonds so these

double and triple bonds are referred

collectively

as multiple bonds

now multiple bonds

are shorter

and a lot stronger

than

conventional single bonds

because

as you

share more and more electrons your

nuclei will get closer and closer

together

which makes them much harder to separate

as well it makes the uh

atomic radius and the

bond length much shorter

now not all

molecules can be represented by these

lewis structures such as the diatomic

fluorine we just studied for example if

you look at

ozone

which is

three oxygen molecules bonded together

uh you may be saying to yourself what's

wrong with this lewis structure well it

can alternately be represented

by

having the single bond on the other side

now the problem with this is that

through experimentation

scientists have found that it doesn't

exist

in one

or two of these states

it exists as sort of an average of the

two so to represent

this property which is called resonance

meaning that

the chemical the chemical is really a

hybrid of two different uh variations of

structure

you take these two lewis structures and

you put an arrow going back and forth in

between them to show that

it's a resonant structure and it can be

shown

in either way but exists in nature

in neither of these two diagrams