Hybridization Theory (English)
Summary
TLDRThis script delves into the significance of hybridization theory in chemistry, essential for understanding molecular geometry and predicting chemical reactivity. It explains how the theory helps visualize molecules in 3D, using carbon as a basis to explore different hybridizations like sp3, sp2, and sp. The script covers the construction of molecules from 2D Lewis structures to 3D models, illustrating concepts with examples like methane, ethylene, and acetylene. It also touches on the impact of lone pairs and steric factors on bond angles, and the importance of hybridization in the broader field of chemistry.
Takeaways
- 📚 Hybridization theory was developed to explain the observed four bonds of carbon and to predict molecular geometry in three dimensions.
- 🧠 Understanding hybridization is essential for visualizing molecules in 3D, which is crucial for grasping chemical and physical properties and predicting reactivity.
- 📐 The theory introduces different hybridization states like sp3, sp2, and sp, which correspond to different bond angles and molecular geometries.
- 🌐 Hybridization allows chemists to understand the orientation of atoms within molecules, which is foundational to molecular geometry.
- 🔬 The carbon atom, with its four valence electrons, can form four covalent bonds through hybridization, contrary to the initial two unpaired electrons in the 2px and 2py orbitals.
- 📈 Hybrid orbitals are created by mixing atomic orbitals, resulting in new orbitals that are used for bonding and have specific spatial arrangements.
- 🛠 The sp3 hybridization results in a tetrahedral arrangement with bond angles of 109.5 degrees, as seen in methane (CH4) and ethane (C2H6).
- 🔬 The sp2 hybridization leads to a trigonal planar arrangement with bond angles of 120 degrees, as found in ethylene (C2H4).
- 🌉 The sp hybridization results in a linear arrangement with bond angles of 180 degrees, as observed in acetylene (C2H2).
- 🔍 The concept of hybridization can be extended to other elements like nitrogen and oxygen, which also follow the principles of electron pair repulsion and orbital hybridization.
- 🔄 Hybridization theory also helps in understanding the existence of geometric isomers and the energy differences between them due to rotation restrictions around double and triple bonds.
Q & A
Why was hybridization theory developed?
-Hybridization theory was developed to better explain the observed four bonds of carbon and to predict overall molecular geometry, including bond angles in three dimensions.
Why is hybridization theory important in chemistry?
-Hybridization theory is important because it allows chemists to envision molecules in three dimensions, which is essential for understanding chemical properties, physical properties, and predicting chemical reactivity.
How does hybridization theory help in visualizing molecules?
-Hybridization theory provides a framework for understanding the orientation of atoms within molecules in three dimensions, which is crucial for predicting chemical behavior and reactivity.
What is the basic electron configuration of a carbon atom?
-A carbon atom has six electrons with the configuration 1s² 2s² 2p², where the valence or outermost electrons are responsible for bond making and breaking.
How does hybridization theory explain the bonding in methane (CH4)?
-In methane, the carbon atom is sp³ hybridized, which means it has four unpaired valence electrons capable of forming covalent bonds with four hydrogen atoms, resulting in a tetrahedral geometry with bond angles of 109.5 degrees.
What are the different types of hybrid orbitals a carbon atom can have?
-A carbon atom can have sp³, sp², and sp hybrid orbitals, each corresponding to different numbers of electron domains (four, three, and two, respectively) and different molecular geometries.
How does the shape of atomic orbitals influence hybridization?
-The shapes of atomic orbitals, such as the spherical 2s and dumbbell-shaped 2p orbitals, determine how they can mix to form hybrid orbitals, which in turn influences the three-dimensional shape of the molecule.
What is the significance of sigma and pi bonds in the context of hybridization?
-Sigma bonds are formed by the overlap of hybrid orbitals along an axis connecting two nuclei, while pi bonds result from the side-to-side overlap of unhybridized p orbitals. Understanding these bonds is crucial for predicting molecular geometry and reactivity.
How does the concept of hybridization apply to other elements besides carbon?
-The principles of hybridization discussed for carbon can also be applied to other elements, such as nitrogen and oxygen, by counting the number of groups (atoms or lone pairs) around the central atom to deduce the type of hybridization.
What is the role of VSEPR theory in understanding molecular geometry?
-Valence Shell Electron Pair Repulsion (VSEPR) theory allows chemists to predict deviations from ideal bond angles based on electron pair repulsions, including the effects of lone pairs and the presence of different groups around a central atom.
Can you provide an example of how hybridization theory helps in understanding the structure of a more complex molecule like tetrodotoxin?
-While the script does not provide specific details about tetrodotoxin, hybridization theory can be used to understand its complex molecular geometry by determining the hybridization states of its constituent atoms and predicting the orientation of its bonds and functional groups in three dimensions.
Outlines
🌐 Understanding Hybridization Theory
Hybridization theory was developed to explain the three-dimensional orientation of atoms within molecules, which is crucial for predicting chemical reactivity. The theory allows chemists to visualize molecules beyond their two-dimensional Lewis structures. It is particularly important for carbon, which typically forms four covalent bonds, contrary to what its two unpaired electrons in the 2px and 2py orbitals would suggest. The concept introduces the idea of hybrid orbitals, such as sp3, sp2, and sp, which help to predict molecular geometry and bond angles. The sp3 hybridization, characterized by one s and three p orbitals, results in a tetrahedral arrangement with bond angles of 109.5 degrees, as seen in methane (CH4) and ethane (C2H6). This understanding is fundamental to molecular geometry and is extended to other elements beyond carbon.
🔍 Exploring sp3 Hybridization and Molecular Geometry
This paragraph delves deeper into sp3 hybridization, highlighting its role in the molecular geometry of compounds like methane and ethane. The sp3 hybridized carbon atom forms four covalent bonds, with the hybrid orbitals arranging themselves to minimize electron repulsion, resulting in a tetrahedral shape. The paragraph also discusses the representation of these molecules in three dimensions, using drawings that distinguish between atoms in the plane and those out of plane. It introduces the concept of sigma bonds, which result from the overlap of hybrid orbitals, and explores the physical properties of ethane, including free rotation around the carbon-carbon single bond and the concept of conformational analysis, which is key to understanding steric strain and the stability of different molecular conformations.
📐 Transitioning to sp2 and sp Hybridization
The script moves on to describe sp2 and sp hybridization, which involve the mixing of different numbers of s and p orbitals to form hybrid orbitals with distinct shapes and bond angles. sp2 hybridization, with one s and two p orbitals, results in a trigonal planar arrangement, while sp hybridization, with one s and one p orbital, leads to a linear geometry. The paragraph explains how these hybridizations allow for the formation of double and triple bonds, as seen in ethylene and acetylene, respectively. It also touches on the concept of pi bonds, which form from the overlap of unhybridized p orbitals, and the implications of these bonds for the stability and reactivity of molecules.
🔬 Visualizing Molecular Structure with Hybridization
This section emphasizes the importance of visualizing molecular structures in three dimensions using hybridization theory. It explains how to deduce hybridization by counting the number of groups around a central atom, including lone pairs. The paragraph illustrates the construction of molecules like propylene, which contains both sp3 and sp2 hybridized carbons, and carbon dioxide, which uses sp and sp2 hybridized building blocks. The discussion on the orientation of hybrid orbitals and the formation of sigma and pi bonds provides a clear understanding of the molecular structure of these compounds.
🌡 Steric Factors and the Impact on Molecular Geometry
The paragraph discusses how steric factors, such as the presence of larger atoms or lone pairs, can affect the rotation around sigma bonds and the stability of molecular conformations. It also explains how deviations from ideal bond angles can occur due to the repulsion between lone pairs and bonding pairs. The Valence Shell Electron Pair Repulsion (VSEPR) theory is introduced as a tool for predicting these deviations, with a focus on the interactions between lone pairs and bonding pairs in molecules like water and ammonia.
♻️ Inversion of Configuration in Nitrogen Hybridization
This section explores the unique property of nitrogen, which can undergo inversion of configuration, demonstrating the dynamic nature of hybridization. The nitrogen in ammonia can rapidly switch between sp3-like and sp2-like hybridizations, a process that occurs billions of times per second. This phenomenon underscores the importance of hybridization theory in understanding the structure and reactivity of molecules, highlighting its central role in the field of chemistry.
🏆 Hybridization Theory as a Cornerstone of Chemistry
The final paragraph of the script reinforces the significance of hybridization theory in chemistry, describing it as the cornerstone of the discipline. It emphasizes the rewards of being able to visualize molecules in three dimensions, including a deeper understanding of chemical and physical properties and the ability to predict chemical reactivity. The script concludes by highlighting the importance of practice in mastering the concepts of hybridization and the ability to mentally manipulate molecular structures.
Mindmap
Keywords
💡Hybridization Theory
💡Three-Dimensional Visualization
💡Carbon Atom
💡Molecular Geometry
💡Valence Electrons
💡Electron Repulsion
💡Sigma Bonds
💡Pi Bonds
💡Hybrid Orbitals
💡Conformational Analysis
💡VSEPR Theory
Highlights
Hybridization theory was developed to explain the four observed bonds of carbon and predict bond angles in three dimensions.
Visualization of molecules in 3D is an essential skill in chemistry, enhancing understanding of chemical and physical properties and predicting reactivity.
Hybridization theory is fundamental to understanding molecular geometry, from simple ethanol to complex tetrodotoxin molecules.
Carbon's electron configuration and the concept of valence electrons are key to understanding hybridization.
Hybridization theory accounts for carbon's ability to form four covalent bonds, contrary to the two unpaired electrons in the 2px and 2py orbitals.
The sp3, sp2, and sp hybrid combinations for carbon are detailed, explaining different molecular geometries.
The shapes of atomic orbitals (spherical for 2s and dumbbell-shaped for 2p) are crucial for understanding hybrid orbital formation.
Hybridization involves mixing atomic orbitals to form hybrid orbitals with specific shapes and orientations.
The sp3 hybridized carbon forms a tetrahedral shape with bond angles of 109.5 degrees, minimizing electron repulsion.
Methane (CH4) and ethane are examples of simple carbon compounds with sp3 hybridized carbons, illustrating ideal bond angles.
Conformational analysis in ethane is possible due to free rotation about the carbon-carbon single bond.
The sp2 hybridization is characterized by one s and two p parts, leading to a planar trigonal shape for molecules like ethylene.
The pi bond in alkenes, formed by unhybridized p orbitals, is essential for understanding geometric isomerism.
The sp hybridization results in a linear shape with 180-degree bond angles, as seen in acetylene.
Hybridization theory can be applied to other elements like nitrogen and oxygen, considering lone pairs as groups.
Counting groups around a central atom helps deduce hybridization, with four, three, or two groups corresponding to sp3, sp2, or sp, respectively.
Hybridization theory allows for the prediction of molecular conformations and the stability of different conformers.
Deviations from ideal bond angles can be predicted using the VSEPR theory, considering lone pair and bonding pair interactions.
Hybridization theory is foundational to chemistry, providing a comprehensive framework for understanding molecular structures and reactivity.
Transcripts
why was hybridization theory developed
why is this theory so important and how
does it allow the chemist to envision a
molecule in three dimensions in this
chapter we will explore some of the
reasons why hybridization theory was
developed a student's ability to take a
two-dimensional molecule off the
blackboard during lecture and fold it
into three dimensions or envision a
molecule in three dimensions from a page
in a book is arguably one of the most
essential skills when learning chemistry
the chemistry student who has the
ability to visualize molecules in three
dimensions is rewarded with a better
understanding of chemical properties
physical properties and most importantly
the ability to predict chemical
reactivity understanding how atoms
within molecules are oriented in three
dimensions requires an understanding of
hybridization theory from simple
molecules such as ethanol to more
complex molecules such as the highly
toxic tetrodotoxin the concepts of
hybridization are the foundation to our
understanding of molecular geometry
however looking at a two dimensional
lewis structure of molecule affords much
information to the scientist for example
the overall connection of atoms within
the molecular formula
after all different structural and
geometric isomers can be imagined from
relatively simple molecular formulas to
introduce the concepts of hybridization
we will first focus all of our examples
on the carbon atom the basic principles
discussed for the carbon atom can also
be applied to other elements which are
explored in later sections a carbon atom
has six electrons in the following
configuration 1s2 2s2 2p2 the relative
energy electron configuration diagram is
another way to visualize where the
electrons are located each arrow in this
diagram represents one of the individual
electrons of carbon however only the
valence or outermost electrons are
responsible for bond making and bond
breaky thus we can ignore the inert
noble core electrons which we can
represent here is the helium element
this abbreviated electron configuration
quickly allows one to ascertain that
there are four valence electrons it
would appear that there are only two
unpaired valence electrons capable of
forming covalent bonds one in the 2px
and one and the 2py orbital
however it is well-known that carbon
forms a total of four covalent bonds to
attain full valency thus all four
valence electrons must be involved in
bonding hybridization theory was
developed in order to better explain the
four observed bonds of carbon in
addition the hybrid model best explains
overall molecular geometry of carbon in
other words bond angles in three
dimensions can be predicted the possible
hybrid combinations for a carbon atom
are sp3 sp2 NSP and are explained in
detail in subsequent sections
let us first start by examining the
shapes of the atomic orbitals for carbon
the 2's atomic orbital is a sphere and
the three 2p atomic orbitals are shaped
like dumbbells oriented along the three
axes 2px 2py and 2pz think of these four
atomic orbitals as three dimensional
shapes where you are most likely to find
an electron 90 percent of the time
[Music]
notice that the electrons have access to
both lobes for the 2p orbitals starting
from the abbreviated electron
configuration for carbon one can imagine
promoting an electron from the 2's
atomic orbital to the unoccupied 2 PZ
atomic orbital although we now have four
unpaired electrons for bonding we still
can't explain the experimentally
observed bond angles for a tetravalent
carbon thus when we mix the 2's atomic
orbital with all three 2p atomic
orbitals we create four new degenerate
energy hybrid orbitals
the shape of the new sp3 hybrid orbital
is best characterized as one part s and
3 parts P as with all orbitals think of
these hybrid orbitals as in
three-dimensional shapes where you can
find the electron ninety percent of the
time the hybridized carbon now possesses
four unpaired valence electrons and is
said to be an sp3 hybridized carbon when
we superimpose all four sp3 hybrid
orbitals on to the carbon atom it
becomes quite cumbersome and confusing
[Music]
thus we simply show how the electrons in
the hybrid orbitals are oriented in
three dimensions
[Music]
the four new hybrid orbitals attempt to
get as far apart from each other as
possible 109.5 degrees think of this as
the orbitals attempting to minimize
repulsion between them thus they are
oriented towards the corners of a
tetrahedron with all angles at 109.5
degrees your instructor will often draw
the sp3 hybridized carbon on the
blackboard as shown the two solid lines
in this drawing are in the plane of the
board the wedge represents the electron
coming out of the plane of the board and
the dashed line represents the electron
going back behind the plane of the board
each of the four sp3 hybrid orbitals
contains one electron capable of forming
a covalent bond the sp3 hybridized
carbon is now capable of forming four
covalent bonds
here X represents any atom with a
valence electron capable of forming a
covalent bond because the electron
density is symmetrically located about
an imaginary line that runs through the
two adjacent nuclei we call these bonds
Sigma bonds an example of a simple
carbon compound with an sp3 hybridized
carbon is methane ch4 the ideal bond
angles are all 109.5 degrees due to all
four equal in size hydrogen atoms
attempting to get as far away from each
other as possible your instructor will
often draw a methane on the blackboard
as shown again the two solid lines in
this drawing are in the plane of the
board
the wedge represents one of the
hydrogen's coming out of the plane of
the board and the dashed line represents
one of the hydrogen's going back behind
the plane of the board
another simple carbon compound that
utilizes sp3 carbons is FA from the two
dimensional Lewis diagram we see that
each carbon has four single bonds thus
both carbons are SP 3 hybridized
starting with 2 sp3 hybridized building
blocks we can start to construct the
molecule in three dimensions by forming
the C C Sigma bond next the six hydrogen
sigma bonds are formed which affords the
final 3-dimensional structure for ethane
the two-dimensional Lewis diagram for
ethane implies that all four bond angles
are 90 degrees
however employing the basic principles
of hybridization theory we see that the
bond angles are all nearly 109.5 degrees
now that the molecular geometry for
ethane in three dimensions has been
determined we can begin to examine some
of ethane interesting physical
properties
for example free rotation may occur
about the carbon-carbon single bond
which allows us to explore simple
conformational analysis confirmations
are different arrangements of atoms due
to these rotations when we place the
electron density around each hydrogen
atom we see that the hydrogen atoms from
adjacent carbons do not touch to make
this diagram easier to view we will
remove the electron density from two of
the hydrogen atoms from the back carbon
even though the hydrogen atoms from
adjacent carbons do not touch
there is torsional strain due to the
electron clouds of the adjacent carbon
hydrogen bonds which impedes the
rotation about the CC bond this gives
rise to the staggered and eclipsed
confirmations for ethane the difference
in relative energy between these two
confirmations is approximately 3
kilocalories per mole it may be easier
to remember that atoms want to be as far
apart from each other as possible think
of it as less crowding when molecules
are viewed down the CC Sigma bond we
call this a Newman projection often you
may see your instructor represent the
Newman projection on the blackboard as
follows
[Music]
[Music]
[Music]
when we replace one of the hydrogen
atoms with an atom that has a larger
atomic radius than hydrogen steric
factors will arise which will increase
the barrier of rotation as the dihedral
angle changes so does the relative
stability of the molecule
similar to the first step of sp3
hybridization one can imagine promoting
an electron from the 2's atomic orbital
to the unoccupied 2 PZ atomic orbital
mixing the 2's atomic orbital with two
of the 2p atomic orbitals creates three
new degenerate energy hybrid orbitals
the shape of the new sp2 hybrid orbital
is best characterized as one part s and
2 parts p as with all orbitals think of
these hybrid orbitals as 3-dimensional
shapes where you can find the electron
90% of the time the sp2 hybridized
carbon now possesses three electrons in
hybrid orbitals and one electron in an
unhybridized 2 PZ orbital when we
superimpose all three sp2 hybrid
orbitals with the unhybridized 2 PZ
orbital on to the carbon atom it becomes
quite cumbersome and confusing
thus we simply show how the electrons
and the hybrid orbitals are oriented in
three dimensions in addition the
electron density of the unhybridized p
orbital is shrunk to a third of its size
for simpler viewing the three new hybrid
orbitals attempt to get as far apart
from each other as possible again think
of this as the hybrid orbitals
attempting to minimize electron
repulsion thus they are oriented 120
degrees apart
your instructor will often draw the sp2
hybridized carbon on the blackboard is
shown again remember that solid lines
are in the plane of the board wedges are
coming out of the plane of the board and
dash lines are going back behind the
plane of the board before we begin to
show how the sp2 hybrid building block
takes part in bonding it is important to
remember that the electron and the
unhybridized 2p orbital as XS 2 both
lobes each of the three sp2 hybrid
orbitals contains one electron capable
of forming a sigma bond
thus the sp2 hybridized carbon is now
capable of forming three sigma bonds and
one PI bond
[Music]
a simple carbon compound that utilizes
sp2 carbons is ethylene from the two
dimensional Lewis diagram we see that
each carbon forms three sigma bonds and
one PI bond thus both carbons are sp2
hybridized starting with 2 SP 2
hybridized building blocks we can begin
to construct the molecule in three
dimensions by forming the C C Sigma bond
next the four carbon hydrogen sigma
bonds are formed which affords the
planar sigma bond framework for ethylene
to form the second bond between the
carbons called the pi bond we can
imagine that the two adjacent parallel
unhybridized to PZ atomic orbitals
overlap when they overlap the two
electrons can be shared allowing each
carbon to attain full valency the two
dimensional Lewis diagram for ethylene
allows the chemist to view the gross
connectivity of the atoms however no
information is conveyed about the PI
bond
when the molecule is represented in
three dimensions we see that half of the
PI bond is above the plane and the other
half of the PI bond is below the plane
an understanding of this electron
density within a PI bond will become
very useful when predicting reactivity
of alkenes
[Music]
[Music]
to gain a better understanding of the PI
bond we should recall the actual shape
of the unhybridized p orbitals when we
envision the actual shape of these
orbitals overlap between the adjacent P
orbitals as possible which allows for
the sharing of these two electrons
however it is very difficult to draw the
molecule this way thus you will often
see the PI bond represented in its
abbreviated form on the right
understanding the PI bond helps us
realize why geometric isomers are
isolobal geometric isomers have the same
gross connectivity but differ only in
how the groups are oriented in space due
to hindered rotation about the doubly
bonded carbons when we draw an imaginary
line along the axis of the double bond
and then compare groups on each carbon
using the cahn-ingold-prelog sequence
rules we can determine if the groups of
priority are on the same side called the
sis isomer often abbreviated Z
alternatively the groups of priority can
be on opposite sides of the imaginary
line called the trans isomer often
abbreviated e or inter conversion of the
isomers to occur we need to have free
rotation about the carbon-carbon double
bond if this were to happen it would
mean that the PI bond would have to
break which requires approximately 70
kilocalories per mole
this will cause each carbon to lose full
valency due to the 2p orbitals no longer
overlapping which will make the alkene
unstable or higher in relative energy
[Music]
thus at room temperature geometric
isomers are a syllable similar to the
first step of sp3 and sp4 decision one
can imagine promoting an electron from
the 2's atomic orbital to the unoccupied
2p z atomic orbital mixing the 2's
atomic orbital with one of the 2p atomic
orbitals causes two new degenerate
energy hybrid orbitals
the shape of the new SP hybrid orbital
is best characterized as one part s and
one part P as with all orbitals think of
these hybrid orbitals as
three-dimensional shapes where you can
find the electron ninety percent of the
time the SP hybridized carbon now
possesses two electrons in hybrid
orbitals and two electrons in the
unhybridized 2p orbitals but when we
superimpose both SP hybrid orbitals with
the unhybridized two PZ and two py
orbitals on to the carbon atom it
becomes quite cumbersome and confusing
thus we simply show how the electrons
and the hybrid orbitals are oriented in
three dimensions in addition the
electron density the unhybridized p
orbitals is shrunk to a third of their
size for simpler viewing the two new
hybrid orbitals attempt to get as far
apart from each other as possible again
think of this as the orbitals attempting
to minimize electron repulsions thus
they are oriented 180 degrees apart
before we begin to show how the SP
hybrid building block takes part in
bonding it is important to remember that
the electrons and the unhybridized 2p
orbitals have access to both lobes your
instructor will often draw the SP
hybridized carbon on the blackboard as
shown again solid lines are in the plane
of the board shaded lobes are coming out
of the plane of the board and dash lines
are going behind the plane of the board
[Music]
each of the two SP hybrid orbitals
contains one electron capable of forming
a sigma bond the SP hybridized carbon is
now capable of forming two sigma bonds
and two pi bonds a simple carbon
compound that utilizes SP carbons as
f-fine or acetylene from the two
dimensional Lewis diagram we see that
each carbon forms two sigma bonds and
two PI bonds thus both carbons are SP
hybridized starting with two SP
hybridized building blocks we can begin
to construct the molecule in three
dimensions by forming the si si Sigma
bond next the two carbon hydrogen sigma
bonds are formed which affords the
linear Sigma bond framework for ethane
to form the second and third bonds
between the carbons called the PI bonds
we can imagine that the two pairs of
adjacent parallel unhybridized 2p atomic
orbitals overlap when they overlap the
four electrons can be shared allowing
each carbon to attain full valency
the two-dimensional Lewis diagram for
f-fine allows the chemist to view the
gross connectivity of the atoms however
no information is conveyed about the PI
bonds when a molecule is represented in
three dimensions we see that half of
each PI bond is above and below a plane
an understanding of this electron
density within these two pi bonds will
become very useful when predicting
reactivities of alkynes to gain a better
understanding of the two PI bonds we
should recall the actual shape of the
unhybridized 2p orbitals when we
envision the actual shape in these
orbitals overlap between the adjacent 2p
orbitals as possible which allows for
the sharing of these four electrons
however it is very difficult to draw the
molecule this way thus you will often
see the PI bonds represented in the
abbreviated form on the right
[Music]
an easy way to deduce hybridizations is
to count groups around a central atom a
group is defined as another atom or a
lone pair when an atom is surrounded by
four three or two groups it will adopt
the sp3 sp2 or SP hybridizations
respectively a helpful way to remember
this is by adding the exponents together
that should equal the number of groups
around the hybridized atom for SP 3 the
exponents add to 4 thus an SP 3 atom has
four groups or SP 2 the exponents had to
3 thus an SP two hybridized atom has
three groups around it these
hybridizations allow the respective
number of groups to be as far apart as
possible again think of it as all groups
attempting to minimize electron
repulsion
[Music]
although we will use the abbreviated
hybridized building block shown here for
subsequent examples it is important to
recall the actual shape of the
unhybridized and hybridized lobes on
carbon for example the unhybridized
lobes were shrunk to a third of their
size and we simply showed how the
electrons and the hybrid orbitals were
oriented in three dimensions so that the
carbon building block does not become
too cumbersome and confusing a simple
carbon compound that utilizes both sp3
and sp4 vines' propylene from the two
dimensional Lewis diagram we see that by
counting groups we can deduce the
hybridization for each carbon atom four
groups employ the sp3 hybridised
building block and three groups employ
the sp2 hybridized building block
starting with one SP 3 and 2 SP 2
hybridized building blocks we can start
to build the molecule in three
dimensions by forming the carbon-carbon
Sigma framework next the six carbon
hydrogen Sigma bonds are formed followed
by the PI bond affording the final
three-dimensional molecule
notice that the methyl group can freely
rotate about the carbon-carbon Sigma
bond while the pi bond affords no
rotation within a sediment are central
atoms that we have not dealt with yet
nitrogen and oxygen
however we employ the same concept for
deducing hybridization simply count
groups on these atoms which means we
also have to count the lone pairs as
groups or groups around nitrogen and
four groups around carbon allows us to
deduce sp3 hybridization while three
groups around the oxygen and carbonyl
carbon allows us to employ sp2
hybridized building blocks
once all the hybrid building blocks have
been deduced we assemble the Sigma bond
framework attach the hydrogen atoms and
form the double bonds as shown
[Music]
again we see that the ch3 group can spin
freely about the carbon-carbon Sigma
bond while the PI bond affords no
rotation in addition the nh-2 group can
spin freely about the carbon nitrogen
Sigma bond now let's look at a compound
that utilizes an SP hybrid building
block carbon dioxide again from the two
dimensional Lewis diagram we see that by
Counting groups we can deduce the
hybridization for each atom two groups
employ the SP hybridized building block
and three groups employ the sp2
hybridized building block for both
oxygen atoms starting with one SP and
sp2 hybridized building blocks we can
start to construct the molecule in three
dimensions by forming both Co Sigma
bonds for both PI bonds to form we need
to rotate the oxygen on the right
forward so that the adjacent
unhybridized 2p orbitals are parallel
thus both PI bonds form affording the
final three-dimensional molecule notice
that the two PI bonds are perpendicular
to each other in addition the lone pairs
on each oxygen atom are perpendicular to
each other
[Music]
as you become comfortable with the
concepts of hybridization you will be
able to fold two-dimensional lewis
structures into three dimensions in your
mind with practice you will also be able
to allow the molecule to undergo
conformational changes in your mind
while predicting the more stable
conformer due to steric interactions and
other effects here we see that the
methyl group prefers to be in the
equatorial position thus one of the
chair confirmations is favored over the
other
as we have seen ideal bond angles are
obtained from the hybrid building blocks
however deviations from ideal bond
angles can and do occur in virtually all
molecules when groups are not equivalent
for example the sp3 hybridized oxygen of
water as a lone pair lone pair
interaction which will cause the two
hydrogen's to become closer than their
ideal bond angle of 109.5 degrees about
104 degrees VSEPR valence shell electron
pair repulsion theory allows the chemist
to make predictions regarding deviations
from ideal bond angles a general trend
that allows predictions from ideal bond
angles is lone pair lone pair
interactions or greater than lone pair
bonding pair interactions which are
greater than bonding pair bonding pair
interactions the sp3 hybridized nitrogen
within the ammonia molecule has three
lone pair bonding pair interactions
which will cause the three hydrogen
atoms to become closer than their ideal
bond angle of 109.5 degrees about 106
degrees
[Music]
an interesting property of nitrogen is
that it has the ability to undergo
inversion of configuration demonstrating
that hybridizations can transform for
this umbrella-like effective happen we
see that the nitrogen hybridization
appears to change from sp3 to an sp2
like nitrogen and back to sp3 current
Berry's predict that there are 200
billion inversions per second for a
molecule of ammonia
if chemistry is considered to be the
central science then hybridization
theory may be considered the cornerstone
of chemistry
after all the student who has the
ability to visualize molecules in three
dimensions is rewarded with a better
understanding of chemical properties
physical properties and most importantly
the ability to predict chemical
reactivity
you
Посмотреть больше похожих видео
VSEPR Theory and Molecular Geometry
3D Structure and Bonding: Crash Course Organic Chemistry #4
VSEPR Theory: Introduction
Class 11 Chemistry | Chemical Bonding and Molecular Str. in 15 Minutes | Rapid Revision by BP Sir
Chapter 1 Part A: Structure and Bonding, acids and bases
Worked examples: Finding the hybridization of atoms in organic molecules | Khan Academy
5.0 / 5 (0 votes)