Chapter 1 Part A: Structure and Bonding, acids and bases
Summary
TLDRThis 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
🌟 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.
📚 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.
🔬 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.
🧬 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.
🌐 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.
Mindmap
Keywords
💡Organic Chemistry
💡Carbon
💡Hybridization
💡Tetrahedral Geometry
💡Covalent Bonds
💡Lewis Structures
💡Valence Electrons
💡Sigma and Pi Bonds
💡Bond Strength and Bond Length
💡Hybrid Orbitals
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
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