6.2 Covalent Bonding and Molecular Compounds
Summary
TLDRThis 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
🧬 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.
⚛️ 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.
🔗 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.
📝 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.
Mindmap
Keywords
💡Covalent Bond
💡Molecule
💡Molecular Formula
💡Diatomic Molecule
💡Potential Energy
💡Bond Energy
💡Bond Length
💡Octet Rule
💡Lewis Structure
💡Resonance
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
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