Chapter 1 Part A: Structure and Bonding, acids and bases

Summary

This chemistry lesson delves into the fundamentals of organic chemistry, exploring the nature of carbon-containing compounds and their prevalence in living organisms. It reviews atomic structure, focusing on electron configurations and orbitals, which are crucial for understanding bonding. The lecture covers various bonding types, including ionic and covalent, with an emphasis on carbon's ability to form four bonds, leading to structures like tetrahedral methane. It also discusses hybridization, including sp3, sp2, and sp, and their roles in forming single, double, and triple bonds. Practical examples like chloroform and acetaldehyde are used to illustrate electron dot structures and line bond diagrams, providing a comprehensive foundation in organic chemistry.

Takeaways

• 🌿 Organic chemistry is the study of carbon-containing compounds and is fundamental to living organisms, including proteins, DNA, food, and medicines.
• 🔬 The modern definition of organic chemistry has evolved from compounds derived from living organisms to encompass the study of all carbon-containing compounds.
• 🌐 Carbon is central to organic chemistry, with 90% of over 30 million chemical compounds containing carbon, often combined with hydrogen, nitrogen, oxygen, halogens, phosphorus, and sulfur.
• ⚛️ Atoms have a positively charged nucleus with protons and neutrons, and negatively charged electrons in orbitals, which can be s, p, or d-orbitals, crucial for understanding chemical bonding.
• 🔑 Carbon atoms are unique, forming four bonds, leading to a tetrahedral geometry that maximizes the distance between bonds, as observed in methane.
• 📚 The script reviews different ways to represent organic compounds, including electron dot structures (Lewis structures), cutaway structures, and Kekulé structures, each showing bonding and electrons differently.
• 🔗 Bonding in organic chemistry primarily involves covalent bonds, which are formed by the sharing of electrons, resulting in a stable electron configuration, in contrast to ionic bonds that involve electron transfer.
• 🧬 Organic molecules can be represented in 3D using various line notations, such as solid, dashed, and wedged lines, to indicate the spatial arrangement of atoms, crucial for understanding molecular geometry.
• 🧪 The script introduces the concept of hybridization, explaining sp3, sp2, and sp hybridizations that influence the shape and bond angles in molecules like methane, ethylene, and acetylene.
• 📈 Bond strength and length are key properties influenced by the type of bond and hybridization, with double and triple bonds being shorter and stronger than single bonds.
• 🧩 Practice problems in the script illustrate how to determine the molecular formula and structure of organic compounds, highlighting the importance of understanding valency and bonding.

Q & A

What is the current definition of organic chemistry?

The current definition of organic chemistry is the study of carbon-containing compounds.

What percentage of chemical compounds contain carbon?

Ninety percent of more than 30 million chemical compounds have carbon in them.

What are the main elements found in organic compounds?

Organic compounds mainly contain carbon, hydrogen, nitrogen, oxygen, and sometimes halogens, phosphorus, and sulfur.

What is the significance of the tetrahedral shape in organic chemistry?

The tetrahedral shape is significant in organic chemistry as it represents the spatial arrangement of atoms bonded to a carbon atom, which typically has four bonds.

How is the structure of methane (CH4) described in terms of hybridization?

The structure of methane (CH4) is described as having sp3 hybridization, where one s orbital and three p orbitals combine to form four equivalent unsymmetrical tetrahedral orbitals.

What is the bond angle in a tetrahedral molecule?

The bond angle in a tetrahedral molecule is approximately 109.5 degrees.

What is the difference between sp2 and sp hybridization?

sp2 hybridization involves one s orbital and two p orbitals, resulting in a planar structure with bond angles of 120 degrees, while sp hybridization involves one s orbital and one p orbital, resulting in a linear structure with bond angles of 180 degrees.

How does the bond strength and bond length change with the type of bond in carbon compounds?

The bond strength increases and the bond length decreases with the number of bonds in carbon compounds. For example, single bonds are weaker and longer than double or triple bonds.

What is the significance of sigma and pi bonds in organic chemistry?

Sigma bonds are the first bonds formed between atoms in covalent bonding, involving head-on overlap of orbitals. Pi bonds are formed by the sideways overlap of p orbitals, and they occur in conjunction with sigma bonds in double and triple bonds.

How are electron dot structures different from line bond structures in representing organic compounds?

Electron dot structures, or Lewis structures, show the valence electrons as dots and indicate covalent bonds by sharing electrons, while line bond structures simplify the representation by showing bonds as lines without the electrons.

Why is it important to understand hybridization when studying organic chemistry?

Understanding hybridization is important in organic chemistry because it helps explain the geometry and reactivity of molecules, as well as the types of bonds that can form between atoms.

Outlines

00:00

🌟 Introduction to Organic Chemistry

The script begins with an overview of organic chemistry, emphasizing its relevance to living organisms and everyday life through examples like proteins, DNA, food, and medicines. It explains the historical definition of organic compounds as those derived from living things and the modern definition focusing on carbon-containing compounds. The importance of carbon in organic chemistry is highlighted, with a discussion of the elements commonly found in organic compounds, such as hydrogen, nitrogen, oxygen, halogens, phosphorus, and sulfur. The atomic structure is reviewed, including the roles of protons, neutrons, and electrons, and the significance of orbitals in chemistry is introduced with descriptions of s, p, and d orbitals. The paragraph concludes with an explanation of how carbon atoms form four bonds, leading to a tetrahedral geometry, using methane as an example.

05:01

📚 Understanding Bonding and 3D Structures

This section delves into the representation of 3D molecular structures using various line conventions, including solid, dashed, and wedged lines. It explains how to interpret and draw these lines to represent bonds that are in the plane of the paper, coming out of the paper, or going into the paper. The script uses dichloromethane (CH2Cl2) as an example to demonstrate how to draw its tetrahedral structure. It also touches on the reasons atoms form bonds, distinguishing between ionic and covalent bonds, and introduces different ways to represent organic compounds, such as electron dot structures, cutaway structures, and Kekulé structures. The importance of achieving a stable electron configuration, typically an octet, is emphasized.

10:02

🔬 Valence Electrons and Bonding

The paragraph discusses the role of valence electrons in bonding, with a focus on how atoms like carbon, nitrogen, oxygen, and halogens form bonds. It provides methods to remember the bonding capacity of these atoms, either by looking at their electron configurations or by using the group number on the periodic table. The paragraph also covers how to draw structures of organic compounds like methane, ethane, butane, and pentane, which illustrate increasing carbon chain lengths. It presents practice problems for drawing electron dot and line bond structures, such as for chloroform (CH3Cl3), and discusses the impossibility of a C3H9 formula due to the saturation of bonding sites on carbon atoms.

15:04

🧬 Covalent Bonding and Hybridization

This section introduces valence bond theory, explaining how covalent bonds form through the overlap of singly occupied orbitals. It describes the formation of a sigma bond between two hydrogen atoms as an example. The theory, developed to describe covalent bonding, is further illustrated with the sp3 hybridization in methane (CH4), where the carbon atom forms four equivalent tetrahedral orbitals. The script discusses bond strength and length, and how they relate to stability, using methane as a model. It also presents different models of methane, including space-filling, Keck whole, and ball and stick representations, to visualize the molecular structure.

20:08

🌐 Hybridization and Bonding Trends

The paragraph explores different types of hybridization, including sp2 and sp, which are responsible for double and triple bonds, respectively. It explains the formation of sigma and pi bonds from sp2 and p orbitals, using ethylene and acetylene as examples. The script also compares the structures, bond strengths, and bond lengths of carbon-carbon and carbon-hydrogen bonds in methane, ethane, ethylene, and acetylene, highlighting trends such as increasing bond strength and decreasing bond length with more bonds. The discussion concludes with the task of drawing electron dot and line bond structures for acetaldehyde (CH3CHO) and acetyl nitrile (CH3CN), emphasizing the types of bonds present and the hybridization of carbon atoms in these molecules.

Keywords

💡Organic Chemistry

Organic chemistry is the study of carbon-containing compounds, which are the building blocks of life. It encompasses the chemistry of living organisms and is central to understanding the molecular basis of biological processes. In the script, organic chemistry is introduced as the chemistry of living things, including proteins, DNA, food, and medicines. Examples given include Vioxx, Lipitor, OxyContin, and benzyl penicillin, highlighting the practical applications of organic chemistry in medicine.

💡Carbon

Carbon is the most important element in organic chemistry due to its unique ability to form four covalent bonds, creating a diverse range of complex molecules. The script emphasizes that 90% of more than 30 million chemical compounds contain carbon, which is essential for the structure and function of organic molecules.

💡Hybridization

Hybridization refers to the concept in which atomic orbitals mix to form new hybrid orbitals, which are used in the bonding of atoms. The script discusses sp3, sp2, and sp hybridizations, which result in different molecular geometries and bond angles. For instance, sp3 hybridization leads to a tetrahedral geometry, as seen in methane (CH4), while sp2 hybridization results in a planar geometry with 120-degree bond angles, as in ethylene.

💡Tetrahedral Geometry

Tetrahedral geometry is a molecular shape where four bonds are arranged around a central atom with bond angles of approximately 109.5 degrees. This geometry is common in organic molecules where a carbon atom is bonded to four other atoms, as explained in the script with the example of methane.

💡Covalent Bonds

Covalent bonds are formed when two atoms share electrons, resulting in a stable molecular structure. The script explains that in organic compounds, electrons are shared equally between atoms, as opposed to ionic bonds where electrons are transferred. Covalent bonds are central to the structure of organic molecules.

💡Lewis Structures

Lewis structures, also known as electron dot structures, are diagrams that represent the valence electrons of atoms as dots and show how these electrons are shared in covalent bonds. The script uses Lewis structures to illustrate how atoms form bonds, such as in acetaldehyde and acetyl nitrile.

💡Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom and are involved in chemical bonding. The number of valence electrons determines the bonding capacity of an atom. The script mentions that carbon has four valence electrons, which it uses to form four bonds, a fundamental concept in organic chemistry.

💡Sigma and Pi Bonds

Sigma (σ) bonds are the strongest type of covalent bond, formed by head-on overlap of atomic orbitals. Pi (π) bonds are formed by the sideways overlap of p orbitals, above and below the plane of the sigma bond. The script explains that double and triple bonds consist of one sigma bond and one or two pi bonds, respectively, and these are found in molecules like ethylene and acetylene.

💡Bond Strength and Bond Length

Bond strength refers to the energy required to break a bond, while bond length is the distance between the nuclei of two bonded atoms. The script notes that bond strength increases and bond length decreases with the number of bonds, which is consistent with the trend observed in carbon-carbon and carbon-hydrogen bonds in various hydrocarbons.

💡Hybrid Orbitals

Hybrid orbitals are orbitals that are formed by the mixing of atomic orbitals. The script discusses sp3, sp2, and sp hybrid orbitals, which are crucial for understanding the geometry and bonding in organic molecules. For example, sp3 hybridization results in a tetrahedral geometry, as seen in methane, where carbon forms four bonds with hydrogen atoms.

Highlights

Introduction to organic chemistry and its relevance to living things, including proteins, DNA, food, and medicines.

Historical definition of organic chemistry as compounds from living organisms has evolved to the study of carbon-containing compounds.

Carbon is the central element in organic chemistry, with over 90% of 30 million chemical compounds containing carbon.

Explanation of the periodic table elements commonly found in organic compounds, such as hydrogen, nitrogen, oxygen, halogens, phosphorus, and sulfur.

Review of atomic structure, including the nucleus, protons, neutrons, and electrons in orbitals.

Description of s, p, and d orbitals and their shapes, with a focus on p orbitals' importance in organic chemistry.

Observation that carbon atoms form four bonds, leading to a tetrahedral geometry in molecules like methane.

Use of solid, wedged, and dashed lines to represent 3D structure in 2D drawings of organic molecules.

Explanation of how atoms form bonds for increased stability, with a distinction between ionic and covalent bonds.

Different ways to represent organic compounds, including electron dot structures, cutaway structures, and Kekulé structures.

Guidelines on how to draw structures of organic compounds, emphasizing the importance of understanding valence electrons.

Practice problems for drawing electron dot and line bond structures, such as for chloroform (CH3Cl).

Explanation of valence bond theory and the formation of covalent bonds through orbital overlap.

Details on bond strength and bond length, and how they relate to the stability of a compound.

Hybridization of orbitals in organic chemistry, including sp3, sp2, and sp hybridizations and their geometric implications.

Comparison of bond strengths and lengths in different types of bonds, such as single, double, and triple bonds.

Drawing the electron dot and line bond structure for acetaldehyde (CH3CHO), including identification of sigma and pi bonds.

Drawing the electron dot structure for acetyl nitrile (CH3CN), highlighting the carbon-nitrogen triple bond.

Transcripts

so in Chapter one we're going to brief a

look at structure and bonding and also

acids and bases so it's going to be a

little bit of a review of general

chemistry and then we go into some acid

and base chemicals what is organic

chemistry organic chemistry is living

things they're all made up of organic

chemicals proteins which makes up hair

and nails DNA is organic chemistry this

controls your genetic makeup foods are

considered organic chemistry and

medicines are also organic chemistry now

you can see lots of different

medications shown to the right here we

have Vioxx lipitor those are cholesterol

medicines oxycontin you may be familiar

with that very addictive drug

cholesterol and benzyl penicillin

previously organic chemistry was defined

in the mid 1700s as a compound that came

from living organisms such as plants and

animals the current definition of

organic chemistry however is just the

study of carbon containing compounds

ninety percent of more than 30 million

chemical compounds have carbon in them

and for the most part the organic

chemicals contain the elements that you

can see colored in the periodic table

they have hydrogen's the carbon here is

the most important one nitrogen oxygen a

few of the halogens and also phosphorous

and sulfur so those compounds combined

with carbon are organic compounds if you

recall from general chemistry the

structure of an atom we have a

positively charged nucleus which is in

the center it's very dense

and contains protons and neutrons and is

also small at 10 to the negative 15

meters the negatively charged electrons

are in the cloud which surround the

nucleus so you can see the nucleus here

protons have a positive charge neutrons

have zero charge and the electrons have

a negative charge so you can see the

negative cloud surrounding the neutron s

and P orbitals are the most important in

organic and biological chemistry S

orbitals are spherical

and have the nucleus at the center so

this one here is an S orbital P orbitals

are dumbbell shaped or kind of look like

peanuts and the nucleus is at the middle

so you can see that nucleus there and

then for the p orbital here and d

orbitals are elongated dumbbell shapes

and the nucleus is at the center there

as well

this one is the d-orbital and there is

the nucleus in the center in green each

of the shells are three perpendicular P

orbitals of equal energy and the lobes

of the p orbital are separated by a

region of no electron density which is

referred to as a node there's no

electron density in here and that is

considered the node when we look at

different orbital diagrams an S shell is

lowest in energy this can hold two

electrons so any of these orbitals can

at most hold two electrons the first

shell which is the 1s holds only two

electrons when we go to the second shell

this has a 2's orbital and also has

three 2p orbitals which can hold each

two electrons as well for a total of

eight electrons and the third shell

which is an S orbital 1s orbital and

then 3p orbitals and we have five D

orbitals that can hold ten electrons for

a total of 18 for the third shell and

these increase in energy as we fill more

shells Cocola and Cooper independently

observed that carbon always has four

bonds the atoms don't have specific

directions but they want to be as far as

possible from each other so if we look

at a carbon this is the structure of

methane this is a carbon with four

hydrogens on it this is what they have

observed and what they notice is that

these hydrogen's want to be as far apart

as possible and this is what it looks

like in 3d shape so this is referred to

as a tetrahedral shape and they get that

from this geometric structure here where

these are as far apart as possible you

might notice that there's lines that are

straight we have some dashed lines and

some wedged lines the two lines that are

just straight are representing bonds in

the page plane

so if you're looking at the lines of

your paper these are in the plane with

your paper the dashed lines are or bonds

that are going away from your paper

sticking out the back of your paper

and then the wedge line refers to a bond

that's

out of the paper

so this would be your tetrahedral

Adam each carbon atom four bonds

when we draw the 3d shape there will be

two bonds in the plane one bond sticking

out and one sticking back

you

so if we're to draw ch2cl2 which is

dichloromethane using solid and wedged

dashed lines to show its tetrahedral

geometry I would start first with just

drawing it flat so carbon we know makes

four bonds and there's four atoms here

so i have h h CL CL so this would be

ch2cl2 this is all flat in your paper if

we want to showing wedges and dashes we

can choose any two things to be flat and

then the other two wedged or - so if I

opt for my hydrogen's to be flat one of

the chlorines must be wedged the other

one must be dashed

you could also draw your chlorines flat

a hydrogen wedged one dashed

one of each of these can be in the plane

and you can drop that

so there's different ways to draw them

and which makes sense because you can

pick up a molecule and rotate it around

any way you like so next you can convert

the following structure into a structure

using wedged normal and dashed lines to

represent the 3d structure so this is

what we're seeing here we have to decide

what is in the plane what's coming out

and what is going away so to me this

looks like this is what is in the plane

because it looks flat so if I were to

draw this the black balls represent

carbon the gray ones are hydrogen so

this is carbon and on this particular

carbon that one looks like it's going

back so I would draw the dashes this one

looks like it's coming forward so

there's a wedge there this one is

pointing back and that one is come

at you so this carbon has two bonds in

the plane one is wedge 2 one is dashed

this carbon has two bonds in the plane

one is wedged one is dashed atoms can

form bonds because the compound that

results is usually more stable than

having the atoms be separate altogether

ionic bonds are those that are insults

they result from electron transfers

you

covalent bonds are bonds that form in

organic compounds and that's resulting

from the sharing of electrons

you

ionic bonds are resulting of the sodium

ion giving all of its electrons to

chlorine and chlorine holds on to them

in an organic compound the electrons are

shared equally between the two atoms we

can represent organic compounds in

different ways

there's the electron dot structures

which you've seen previously are also

known as Lewis dot structures so Lewis

structures our electron dot structures

which show the valence electrons of an

atom as dots the valence electron are

shown as dots as we've seen before and

notice also that oxygen and nitrogen

which have extra lone pairs those

electrons are shown as well

the cutaway structures have the line

drawn between the two atoms which

indicates that there is a covalent bond

there represented by the sharing of two

electrons in a stable molecule the atom

will have eight electrons which is a

completed shell or four hydrogen there's

only two electrons there each of these

atoms that are not hydrogen will have a

complete octet these show all of the

individual electrons you can see carbon

has the four electrons hydrogen

contributes one and those are each

forming a bond there's also the kekulé a

structures which just shows the bond as

a line instead of two individual

electrons so this is what you can see

for a carbon this is nitrogen oxygen and

a carbon that's bonded to an oxygen

carbon here you notice is making four

bonds nitrogen makes only three bonds

oxygen makes two bonds here carbon has

four carbon can form four bonds still

one two oxygen this oxygen has two bonds

when we look at these atoms these are

the main ones that you'll need to know

for again at chemistry we don't need to

be concerned about electron dot

structures for the rest of the periodic

table hydrogen has one bond carbon makes

four bonds

nitrogen has three bonds oxygen two

bonds and the halogens have one fun

available if you have a hard time

remembering how many bonds each atom mix

you can look at the electron

configuration so for carbon it has 1s2

2s2 2p2 this first shell is already full

there's no electrons available for

bonding here but with two s 2 and 2 P 2

these have four valence electrons still

available for bonding four valence and

each of those is available for a bonding

like I said and that's how we can

remember how many there are you could

also look at the group number of the

group number here is four five six and

that is seven which tells you how many

electrons there are and as you put them

around the atom this would be 1 2 3 4 5

2 of these are lone pairs and that's why

we have free bonds here this is 1 2 3 4

5 6 and these have 7 these are all

paired ups there's one bond available

you

those are the two tricks you can use to

remember how many bonds each of those

atoms make for drawing structures of

organic compounds methane is one that

has two carbons this is C C there's

enough hydrogen's around here to fill in

the rest of the remaining bonds this

ends up being C 2 H 6 propane has 3

carbons and each of these carbons has 4

bonds each of the hydrogen has one bond

butane has 4 carbons

you

and pentane has five

you

here are a couple of practice problems

you can pause the video if you want and

try to work out the problems on your own

before you come back to look at the

answers we're looking at drawing both

electron dot and line bond structures

for chloroform chloroform is ch3cl 3 you

can start with carbon which makes has 4

valence electrons hydrogen has 1 and

chlorine has 7 and we have three of

those one of these can form a bond

together one of these can and then the

other chlorines can as well

my resulting electron dot structure

you

would look like this and the line bond

would look like this in the line bond

drawing you don't need to include all of

the electron lone pairs for chlorine if

we look at the next problem it's asking

what is select formula for the following

and we can figure this out based on the

number of bonds that carbon needs to

make carbon has four valence electrons

so it needs to make four bonds this is

only bonding to chlorine so this one

needs to be four nitrogen is in group

five it has five electrons two of these

are already paired up so it's available

to make three bonds hydrogen is the only

thing it's bonded to so this should be 3

4 CH blank o H it is and carbon makes

four bonds oxygen has six one two three

four five six so there's two bonds

available hydrogen does one so if I look

at this carbon it has a bond to H oh

this is to an H so this one's happy so

we need to fit in to more H's here in

order to fill that so this carbon has

four bonds this oxygen has two so why

can't inorganic molecule have the

formula C 3 H 9 if we DRI tried to draw

this out

I have three C's this has four bonds

this has four bonds I have one two three

four five six seven eight hydrogen's all

of my carbons have four bonds and I

can't fit any more on there there's not

enough bonds available if we look at

valence bond Theory covalent bonds form

when two atoms come close to each other

so that is singly occupied orbital

another one overlaps we have two

individual s orbitals and these

hydrogen's they come close together

where their orbitals can overlap and

they can form in h2 molecule a valence

bond theory was developed in order to

model and describe the covalent bonding

in the valence bond three electrons are

paired in overlapping orbitals and are

attracted to nuclei of both atoms so I'm

going to use the e minus here to

represent electrons

you

the electrons are attracted to nuclei of

both atoms a hydrogen bond can results

from the overlap of two singly occupied

hydrogen s orbitals like we said here

these two singly occupied ones so when

two of these s orbitals get together to

form a new one it forms a sigma bond so

if we had H H got together it forms H H

this is a sigma bond so two s orbitals

form a sigma bond

you

and that is a face on or head on bond

bond strength is the energy that's

released when a bond forms

and the bond length is the optimum

distance between a nuclei leading to

maximum stability

you

so there's going to be an optimum bond

length versus bond strength this is too

close together this is an appropriate

bond length and this is too far apart if

we look at the sp3 orbital and structure

of methane which is ch4 a carbon has

four valence electrons like I mentioned

that 2 SP 2 and 2 P 2 in ch4 all of

these are identical or tetrahedral so an

sp3 hybrid orbital has 1s orbital and 3p

orbitals and those will combine to

perform four equivalent unsymmetrical

tetrahedral orbitals

you

so there's one s and three P's which

make the SP three this was figured out

by Pauling in 1931

so sp3 orbitals have a carbon overlap

with 1s orbital and for H orbitals that

are identical so each CH bond has a

strength of 439 kilojoules per mole and

a length of 109 picometers this is the

model of ch4 this is a space-filling

this is the Keck whole a structure and

then we have a ball and stick one here

so this carbon has bonds to equal that

are equal for each of these CH bonds we

can see the two lines in the plane and

then the wedge and the dash here the

bond lengths of each of these is 109

picometers

and the bond angle between these two

here is 109.5 so the bond angle for each

h CH is 109.5 this is the tetrahedral

angle

we have two carbons that are bonded

together by overlap of an sp3 orbital

from each carbon that is what we see in

ethane so a CH bond in ethane is 421

kilojoules per mole

while the carbon-carbon bond is 154

picometers line and it's slightly less

strong at 377 all of the bond angles of

ethane are tetrahedral so this bond here

is 109 point 5 so we have the two carbon

sp3 atoms come together to combine to

form the sp3 sp3 Sigma bond and then we

have the ethane here in the calculate

structure and then the ball and stick

figure so you can see the different bond

lengths here other kinds of

hybridization that we will encounter

include sp2 and SP an sp2 orbital comes

from an sp2 carbon as we can see here

these are made up of sp2 hybrid orbitals

have 1 2 s orbital

that combines with two 2p orbitals

which gives it reor Battelle's

and those are s P P or s P 2 which

results in a double bond so you can see

this double bond here we have an sp2

carbon the green ones are sp2 orbitals

the green orbitals that you see are sp2

orbitals and then we have one p orbital

you

when the p-orbitals combine these get

together this forms a pi bond

and this is also the other part of the

PI bond this is where the electrons are

this is the bonding part this is the

nonbonding part and where the carbons

forms that first bond together that is a

sigma bond so sp2 orbitals are planar

which means they're flat

and they have bond angles of 120 degrees

there

so the sp2 orbitals are planar and flat

the remaini p orbital is perpendicular

to the plane

you

so we get to sp2 hybridized orbitals

with a head-on overlap form a sigma bond

and then we have P orbitals that have

overlap side to side giving us a PI bond

so if you think of to pregnant ladies

with their bellies sticking out if they

were trying to give each other a hug

head-on you'd have the two pregnant

bellies touching each other that would

be your head-on Sigma bond and if they

were then going to try to shake hands

while their bellies were touching that

would be your side to side PI bond click

an sp2 bond is double bond the SP bond

is a triple bond it has a carbon with a

2's orbital that hybridizes with the

single p orbital which gives to SP

hybrids so two of these P orbitals are

unchanged these are linear and 180

degrees apart on an x-axis so 2p

orbitals are perpendicular on the y and

z x axis so this would be one SP orbital

one SP hybrid this is the other and they

will combine here to make the triple

bond which you'll see in the next slide

so we have two SP orbitals that form

this SP head-on and that forms this

Sigma bond here the two P orbitals here

are going to form that PI bond so the SP

orbitals are responsible for the Sigma

bond the P orbitals are responsible for

the PI bond so two SP hybrid orbitals

from each carbon form the SP SP Sigma

bond so Sigma bond comes from SP SP

orbitals and the PI bond comes from the

P orbitals

just like before the SPR going to have

that head-on overlap and then the

p-orbitals will have the sideways

overlap and you can see the space

filling picture here you can see that

that is 180 degrees this is linear and

that's what the structure of c2h2

so if we compare the carbon carbon and

carbon hydrogen bonds and methane ethane

ethylene and acetylene you can see the

different structures here the bond

strength increases as you increase the

number of bonds and the bond length

decreases

you

and you really want to pay attention to

the carbon carbon the trends are the

same for hydrogen but if we're comparing

carbon carbon bonds here the 377 728 965

these bonds are getting stronger if we

look at kilojoules per mole and the bond

length goes from 154 134 to 120 are

getting shorter the bond strength

increases with the number of bonds and

the Monst length decreases with the

number of bonds and that's because the

electrons are being held closer together

you

you

and the CH bonds have the same trend

where they increase and the same with e

bond lengths okay so we are tasked with

drawing the electron dot and line bond

structure for acetaldehyde or ch3cho so

we've want to draw this lewis structure

or the calculate structure first so i

have c h and sometimes when organic

compounds are written out the order in

which it's written kind of is an

indicator as to how things are put

together so I have a C with three ages

bonded to a C an H and an O so I'm going

to look at this and notice that this

carbon has four bonds this one only has

three oxygen is only having one oxygen

has six valence one of them is used here

carbon has one left over so what I'll do

is join these two electrons together

and make a double bond there okay so

that is the electron dot structure for

acetaldehyde so what types of bonds are

present so all of these bonds here the

CHS and the c2c those are Sigma bonds

and that is the CH the c2c those are the

Sigma bonds and I see I have a double

bond here which means that this carbon

is sp2 and this is forming a PI bond

that's the C to O is a PI bond what does

the hybridization of each carbon this

one here with all the single bonds that

is an sp3 this one here is sp2 the bond

angles for sp3 our 109.5 for sp2 are 120

so next we want to draw

electron dot structure for acetyl

nitrile which is ch3 CN and then that

tells us that we can have a carbon

nitrogen triple bond so again if we look

at how this is written out I have a C

with three H's to a C to an N carbon has

four electrons nitrogen has five and

we're told that the CN has a triple bond

so this is going to become those

electrons come here these come here

and then I have electron lone pair left

over and that's how we would draw the

triple bond and this structure

you