Sigma and Pi Bonds (A-Level Chemistry)
Summary
TLDRThis informative video from chemistrystudent.com dives into the intricacies of sigma and pi bonds, essential concepts in understanding covalent bonding. It explains that covalent bonds form when two atoms share a pair of electrons, with the bond's strength and properties depending on the type of overlap between atomic orbitals. Sigma bonds result from a direct overlap, positioning the bonding orbital close to both nuclei for a strong, non-rotatable bond. In contrast, pi bonds arise from the sideways overlap of p orbitals, creating a weaker, less stable bond with restricted rotation. The video uses ethane and ethene as models to illustrate single and double bonds, respectively, highlighting the increased reactivity of double-bonded molecules like ethene due to the presence of pi bonds. The content is enriched with visual explanations and is aimed at viewers with a foundational understanding of chemistry.
Takeaways
- 🔬 Covalent bonds are formed when two atoms share a pair of electrons, resulting from the overlap of half-filled atomic orbitals from each atom.
- 🧲 The attraction between the positively charged nuclei and the shared electrons in the new orbital pulls the atoms together, forming a covalent bond.
- 📍 Sigma bonds are created from the direct overlap of atomic orbitals, resulting in a strong bond that is close to the nuclei of both atoms.
- 📏 Pi bonds form from the sideways overlap of p-shaped orbitals, creating a weaker bond that is further from the atomic nuclei compared to sigma bonds.
- 🚫 Pi bonds restrict rotation due to the overlapping p orbitals, which can twist and break if rotated too far.
- 🔄 The formation of a sigma bond from one pair of half-filled p orbitals leaves the remaining p orbitals free to potentially form a pi bond.
- 🛑 Sigma bonds are always stronger than pi bonds, which contributes to the differing reactivity between molecules with single and double bonds.
- ⚙️ In ethane, carbon atoms are connected by a sigma bond, with each carbon atom having three bonds to hydrogen atoms.
- ⚓️ In ethene, a carbon-carbon double bond consists of one sigma bond and one pi bond, represented by two lines between the carbon atoms.
- 🔗 A double bond is stronger and less likely to break than a single bond, but a carbon-carbon double bond is not twice as strong as a single bond due to the weaker pi bond.
- 🌟 Alkenes, like ethene, are more reactive than alkanes, such as ethane, because they contain pi bonds that are easier to break, leading to increased reactivity.
Q & A
What is a covalent bond?
-A covalent bond is a type of atomic bond formed when two atoms share a pair of electrons. It occurs when half-filled atomic orbitals from two different atoms overlap, creating a bonding orbital where the pair of electrons can exist, attracting both positively charged nuclei to the electron density and pulling the atoms together.
What are the different shapes of atomic orbitals that can be involved in covalent bonding?
-The different shapes of atomic orbitals that can be involved in covalent bonding include s and p shaped orbitals, which are the most commonly studied. Other half-filled atomic orbitals can also overlap or merge to create a bonding orbital between two atoms.
How does a sigma bond form?
-A sigma bond forms when two atomic orbitals face each other and overlap easily. The resulting bonding orbital is close to the nuclei of both atoms, with both nuclei having a high level of attraction to the electrons in the orbital, making the bond very strong.
What is a pi bond and how does it differ from a sigma bond?
-A pi bond is a covalent bond that forms from the sideways overlap or bending of p-shaped atomic orbitals from two atoms. Electrons in a pi bond are further from the nuclei of both atoms compared to a sigma bond, resulting in a weaker attraction and a bond that is easier to break. Unlike sigma bonds, pi bonds have restricted rotation due to the two areas of electron density above and below the sigma bond.
How does the bonding in ethane (C2H6) differ from that in ethene (C2H4)?
-In ethane, carbon atoms are bonded by a sigma bond formed by the direct overlap of two orbitals, with each carbon atom having three bonds to hydrogen atoms. In ethene, the carbon atoms are bonded by both a sigma bond and a pi bond, forming a carbon-carbon double bond. Each carbon atom in ethene has two bonds to hydrogen atoms, with two half-filled p-shaped orbitals left over for the pi bond formation.
Why are double bonds stronger than single bonds?
-Double bonds are stronger than single bonds because they involve both a sigma bond and a pi bond between the atoms. The presence of two bonds results in a higher bond energy and a stronger attraction between the atoms, making the double bond harder to break than a single bond.
Why are alkenes more reactive than alkanes?
-Alkenes are more reactive than alkanes because they contain a carbon-carbon double bond, which includes a pi bond that is weaker and more easily broken than a sigma bond. This higher reactivity allows alkenes to undergo a wider range of chemical reactions compared to alkanes.
What happens if a pi bond is broken?
-If a pi bond is broken, the sigma bond between the atoms still remains intact. This leaves the molecule with a single bond between the previously double-bonded atoms, which can lead to different chemical reactivity and properties compared to when the pi bond was intact.
What is the significance of the orientation of p orbitals in forming sigma and pi bonds?
-The orientation of p orbitals is significant because it determines the type of bond that can form. When p orbitals overlap directly, they form a sigma bond. However, when they cannot overlap directly due to their orientation, they can bend and overlap sideways to form a pi bond. The three p orbitals (px, py, pz) are oriented at 90 degrees to each other, which allows for the formation of pi bonds after a sigma bond has been established.
How does the rotation of atoms affect the pi bond?
-The rotation of atoms affects the pi bond because the p orbitals that form the pi bond overlap sideways. If one atom rotates, the pi bond's orbitals will also attempt to rotate, which can lead to twisting. If twisted too far, the pi bond can break due to the strain caused by the misalignment of the p orbitals.
What is the role of electron density in the strength of sigma and pi bonds?
-Electron density plays a crucial role in the strength of sigma and pi bonds. In sigma bonds, the electron density is closer to the nuclei of both atoms, resulting in a strong attraction and a robust bond. In pi bonds, the electron density is further from the nuclei, leading to a weaker attraction and a less stable bond compared to sigma bonds.
Why are sigma bonds free to rotate?
-Sigma bonds are free to rotate because the overlapping orbitals that form the bond are directly aligned with the nuclei of the bonding atoms. This direct alignment allows for rotation without disrupting the orbital overlap, which means the bond's strength and integrity are maintained during rotation.
Outlines
🔬 Understanding Sigma and Pi Bonds
This paragraph introduces the concept of sigma and pi bonds, fundamental to understanding covalent bonding in chemistry. It explains that a covalent bond is formed when two atoms share a pair of electrons, with the bonding occurring due to the overlap of half-filled atomic orbitals. The paragraph emphasizes the importance of being familiar with the shapes of s and p orbitals, which are most commonly studied. It also clarifies that double bonds are formed when additional orbitals overlap or merge, and highlights the difference in strength between the first and second bonds in a double bond. The explanation is set to compare sigma and pi bonds by examining the carbon bonding in ethane and ethene.
🌐 Sigma and Pi Bonds in Ethane and Ethene
This paragraph delves into the specifics of sigma and pi bonding by using ethane and ethene as models. It describes how in ethane, carbon atoms are connected by a sigma bond, formed by the direct overlap of orbitals, with each carbon also bonded to three hydrogen atoms. The leftover half-filled p orbitals are noted but do not participate in bonding in ethane. In contrast, ethene features a carbon-carbon double bond, consisting of one sigma bond and one pi bond. The sigma bond is formed by the direct overlap of p orbitals, while the pi bond results from a sideways overlap of the remaining p orbitals. The paragraph explains that the pi bond is weaker and less stable than the sigma bond, which contributes to the higher reactivity of alkenes like ethene compared to alkanes like ethane. It also mentions the restricted rotation in double bonds due to the presence of the pi bond, and concludes by summarizing the key differences between sigma and pi bonds in terms of strength, electron proximity to the nuclei, and rotational freedom.
Mindmap
Keywords
💡Covalent Bond
💡Sigma Bond
💡Pi Bond
💡Atomic Orbitals
💡Ethane
💡Ethene
💡Bonding Orbital
💡Electron Density
💡Reactivity
💡Overlap of Orbitals
💡Double Bond
Highlights
Covalent bonds are formed when two atoms share a pair of electrons, creating a bond through the overlap of atomic orbitals.
Sigma and pi bonds are two types of covalent bonds, differing in the way atomic orbitals overlap.
Sigma bonds result from the direct overlap of orbitals, leading to a strong bond that is close to the nuclei of both atoms.
Pi bonds are formed by the sideways overlap of p-shaped orbitals, resulting in a weaker bond that is further from the atomic nuclei.
Ethane serves as a model for sigma bonding, with carbon atoms bonded by direct orbital overlap.
Ethene (ethylene) demonstrates both sigma and pi bonding, with a double bond consisting of one sigma and one pi bond.
A carbon-carbon double bond in ethene is stronger and less likely to rotate freely compared to a single bond in ethane.
Sigma bonds are represented by a single line, while pi bonds are represented by a double line between carbon atoms.
Although a carbon-carbon double bond is stronger, it is not twice as strong as a single bond due to the relative weakness of the pi bond.
Ethene's higher reactivity compared to ethane is attributed to the presence of the pi bond, which is easier to break.
The orientation of p orbitals in ethene allows for the formation of a pi bond after the sigma bond is established.
The inability to freely rotate a pi bond leads to restricted rotation in molecules with double bonds.
The video explains the concept of atomic orbitals and their role in the formation of covalent bonds.
Different shaped orbitals, such as s and p orbitals, are studied for their role in covalent bonding.
The video uses the example of oxygen molecules to illustrate the formation of sigma bonds through p orbital overlap.
The video provides a detailed comparison between sigma and pi bonds, emphasizing their structural differences and implications for molecular stability.
The video explains that while pi bonds are weaker, their presence in molecules like ethene contributes to their reactivity.
The video concludes by summarizing the key differences between sigma and pi bonds, and their impact on molecular properties.
Transcripts
hello matt here from
chemistrystudent.com in this video we're
going to look at sigma and pi bonds
we're going to talk about what sigma and
pi bonding actually is
how sigma and pi bonds arise and compare
them both by looking at the carbon
bonding in ethane and ephene
covalent bonding has been covered in a
separate video check the links in the
description below
before we talk in detail about sigma and
pi bonding there are a few essential
ideas you need to be comfortable with
a covalent bond is formed when two atoms
share a pair of electrons it is an
example of an atomic bond
when a covalent bond forms half-filled
atomic orbitals from two different atoms
overlap creating a bonding orbital that
a pair of electrons can exist in
electrons are negatively charged and the
positively charged nuclei of both atoms
are attracted to the electron density in
the new in orbital this pulls both atoms
together and creates a covalent bond
this is drawn as a single line between
the two atoms
there are different shaped orbitals
electrons can be in depending on their
distance from the nucleus of an atom
the orbital shape just refers to the
area that an electron pair is likely to
be in at any one time
they are essentially constantly moving
around within this space
at this level s and p shaped orbitals
are the ones most commonly studied
other half-filled atomic orbitals can
sometimes also overlap or merge creating
another bonding orbital between the two
atoms leading to a double bond
the atoms are harder to split apart when
double bonded together although the
second bond is weaker than the first
single bond
recap done let's go
as mentioned to form a covalent bond
between two atoms atomic orbitals must
overlap or merge to create a bonding
orbital
the two atomic orbitals face each other
and if the orbitals can overlap easily
the bond formed is called a sigma bond
shown with the greek symbol for sigma
the bonding orbital is close to the
nuclei of both atoms and the nuclei both
have a high level of attraction to the
electrons in the orbital making the bond
very strong
single bonds are always examples of
sigma bonds as atoms will always try and
arrange themselves to maximize the
overlap of two orbitals
for example in an oxygen molecule each
oxygen atom has a half-filled p-shaped
orbital
the p-shaped orbitals have lobes that
stick out of each side of the nucleus
and the load from each p orbital from
each atom can point towards each other
and the two can overlap and merge
the bond formed is a sigma bond as the
both orbitals overlap directly
the bond is also free to rotate as
rotating each nuclei has no impact on
the bonding orbital
before the oxygen atoms formed a sigma
bond they each had two half-filled p
orbitals the sigma bond was formed from
the direct overlap of one of these from
each atom
this means that each oxygen atom still
has a half-filled p orbital and wants to
form another bond
now things can get a bit interesting
remember that there are three p shaped
orbitals around the nucleus p x p y and
p z
and their lobes are each pointing in
different directions or 90 degrees to
each other
if two half-filled p orbitals overlap
fully to form a sigma bond the
half-filled orbitals left over cannot
directly line up or overlap
they are simply pointing in the wrong
direction
as the atoms are now quite close
together due to the single sigma bond
the other half-filled orbitals can
actually bend inwards slightly and
overlap sideways to create a new bonding
orbital between the two atoms
the electrons in this orbital are
further from the nuclei of both atoms
compared to the sigma bond and the bond
is therefore weaker
this type of bond is called a pi bond
shown with the greek symbol pi
due to each orbital having two lobes the
p orbitals overlap sideways creating two
areas of electron density above and
below the sigma bond already formed
this means the bond is now unable to
rotate freely
if one atom rotates the pi bond in
orbitals will try to rotate as well
meaning they'll become twisted and if
twisted too far will break
sigma bonds are covalent bonds that form
from the direct overlap of two orbitals
from two atoms
electrons in the bond are close to the
nuclei of both atoms and the bond is
strong
pi bonds are covalent bonds that form
from the sideways overlap of two
p-shaped orbitals from two atoms
electrons in the bond are further from
the nuclei of both atoms and as a result
pi bonds are weaker than sigma bonds
at this level the carbon bonding in
ethane and ethene is often used to model
sigma and pi bonding
in ethane the carbon atoms are bonded by
the direct overlap of two orbitals a
sigma bond
each carbon atom has three bonds to
hydrogen atoms and this leaves one
half-filled p-shaped orbital left over
the carbon atoms arrange themselves in
such a way that enables the p shaped
orbitals to overlap
in ifene the carbon atoms are again
bonded by the direct overlap of two
orbitals a sigma bond as well as a
sideways overlap of two other orbitals a
pi bond
this is referred to as a carbon-carbon
double bond
each carbon atom has two bonds to
hydrogen atoms and this leaves two
half-filled p-shaped orbitals left over
to start with the carbon atoms will
arrange themselves to have that direct
overlap of two orbitals and the sigma
bond will form
this will leave two p shaped orbitals
left over that can overlap sideways
between the two atoms forming a pi bond
there are now essentially two bonds
between the carbon atoms a sigma bond
and a pi bond this is called a double
bond and is shown as two lines between
the carbon atoms
one line represents the sigma bond and
the other the pi bond
as the carbon atoms are held by two
bonds a carbon-carbon double bond is
stronger and harder to break than a
single carbon bond
sigma bonds are stronger than pi bonds
however meaning that a carbon-carbon
double bond isn't twice as strong as a
carbon-carbon single bond
it is easier to break a pi bond than a
sigma bond this actually gives ethene a
higher reactivity of an ethane and
explains why alkenes are more reactive
than alkanes
if a pi bond is broken the sigma bond
between the atoms still remains
so to summarize
covalent bonds are formed from the
overlap of two atomic orbitals from two
atoms
sigma bonds form from the direct overlap
of atomic orbitals from two atoms
electrons in the bond are close to the
nuclei of both atoms given a high level
of attraction that is hard to break
making sigma bonds strong
sigma bonds are free to rotate
pi bonds formed from the sideways
overlap or bending of p shaped atomic
orbitals from two atoms
electrons in the bond are further from
the nuclei of both atoms compared to a
sigma bond given a weaker attraction and
a bond that is easier to break
as p shaped orbitals have two lobes a pi
bond has two areas of electron density
one above and one below a sigma bond
already formed between the atoms
there is restricted rotation and the pi
bond is unable to freely rotate and
twist
i hope you found this video useful
please check out other relevant videos
in the links given in the description
below and visit chemistryshooting.com
for free notes and revision materials
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