Hybridization Theory (English)

Summary

This 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

00:00

🌐 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.

05:02

🔍 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.

10:04

📐 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.

15:08

🔬 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.

20:09

🌡 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.

25:14

♻️ 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.

30:16

🏆 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.

Keywords

💡Hybridization Theory

Hybridization Theory is a fundamental concept in chemistry that explains the geometry of molecules. It was developed to account for the observed bonding patterns, particularly the four bonds formed by carbon atoms, which cannot be explained by the traditional electron configuration alone. The theory allows chemists to predict molecular geometry and bond angles in three dimensions. In the video, hybridization theory is essential for understanding how atoms within molecules are oriented and how this orientation affects chemical properties and reactivity.

💡Three-Dimensional Visualization

Three-Dimensional Visualization is the ability to mentally construct and manipulate the structure of molecules in 3D space. This skill is crucial for chemistry students as it enhances understanding of molecular geometry, chemical properties, and reactivity. The video emphasizes the importance of this skill by illustrating how students can translate 2D Lewis structures into 3D models, a process that relies heavily on hybridization theory.

💡Carbon Atom

The Carbon Atom is central to organic chemistry and is the focus of the hybridization theory discussion in the video. With an electron configuration of 1s² 2s² 2p², carbon has the unique ability to form four covalent bonds, which is a key aspect of hybridization theory. The video explains how carbon's valence electrons are involved in hybridization to form sp³, sp², and sp hybrid orbitals, which dictate the geometry of molecules like methane, ethylene, and acetylene.

💡Molecular Geometry

Molecular Geometry refers to the three-dimensional arrangement of atoms within a molecule. It is a direct consequence of hybridization and is crucial for predicting chemical behavior. The video discusses how different hybridizations (sp³, sp², sp) lead to distinct molecular geometries, such as tetrahedral, trigonal planar, and linear, respectively.

💡Valence Electrons

Valence Electrons are the outermost electrons of an atom that participate in chemical bonding. In the context of hybridization theory, the valence electrons of carbon are particularly important as they are involved in the formation of hybrid orbitals. The video explains how carbon's four valence electrons are utilized in forming covalent bonds, which is central to understanding molecular structure.

💡Electron Repulsion

Electron Repulsion is the principle that electrons with the same charge repel each other. In molecular geometry, this principle leads to the arrangement of hybrid orbitals that minimizes repulsion, resulting in specific bond angles. The video uses this concept to explain why sp³ hybrid orbitals arrange themselves in a tetrahedral shape with bond angles of 109.5 degrees.

💡Sigma Bonds

Sigma Bonds are the first type of covalent bond formed between atoms, where the electron density is most concentrated along the axis connecting the nuclei. The video discusses sigma bonds in the context of hybridization, explaining how hybrid orbitals participate in the formation of these bonds, which are a fundamental part of molecular structure.

💡Pi Bonds

Pi Bonds are a type of covalent bond formed by the sideways overlap of p orbitals, resulting in electron density above and below the plane of the bonded atoms. The video explains the formation of pi bonds in molecules like ethylene and acetylene, which are crucial for understanding the double and triple bond structures in these molecules.

💡Hybrid Orbitals

Hybrid Orbitals are a combination of atomic orbitals that result in new orbitals with different energies and shapes, allowing for the formation of the observed number of bonds in molecules. The video describes the creation of sp³, sp², and sp hybrid orbitals from carbon's atomic orbitals, which are essential for understanding the bonding capacity and geometry of carbon-containing molecules.

💡Conformational Analysis

Conformational Analysis is the study of different spatial arrangements of atoms in a molecule, specifically those that can be interconverted by rotation around single bonds. The video uses ethane as an example to illustrate how conformational changes, such as staggered and eclipsed forms, can be understood through the lens of hybridization theory and steric factors.

💡VSEPR Theory

VSEPR (Valence Shell Electron Pair Repulsion) Theory is a model used to predict the geometry of individual molecules based on the repulsion between electron pairs in the valence shell of an atom. The video briefly mentions VSEPR in the context of deviations from ideal bond angles caused by lone pairs or different electron pair interactions, which is important for understanding the actual shape of molecules.

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

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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

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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