CHEM 257 - Fall 2024 - Lecture 1 - Video 2
Summary
TLDRThis chemistry lecture delves into molecular bonding, focusing on ionic and covalent bonds. Ionic bonding involves the transfer of electrons between a metal and non-metal, forming charged ions. Covalent bonding is further divided into polar and non-polar, determined by electronegativity differences and molecular geometry. The lecture also revisits Lewis structures, a fundamental tool for visualizing molecular connectivity, and introduces advanced concepts like multiple bonds and formal charge. The goal is to understand and predict molecular behavior through these bonding principles.
Takeaways
- đŹ Ionic bonding involves a transfer of electrons between a metal (losing electrons, forming cations) and a non-metal (gaining electrons, forming anions), forming compounds like potassium chloride (KCl) and sodium chloride (NaCl).
- đ Polar covalent bonds are based on the difference in electronegativity between atoms, leading to an unequal sharing of electrons and creating a dipole moment, as seen in molecules like HCl.
- đ Non-polar covalent bonds occur when electrons are shared equally between atoms with a small difference in electronegativity, resulting in no net dipole moment, exemplified by methane (CH4).
- đ Molecular geometry and the interaction of lone pairs play a crucial role in determining the polarity of molecules, with symmetry often cancelling out dipole moments, as in the case of carbon dioxide (CO2).
- đ Electronegativity is a key concept in understanding the nature of chemical bonds, with the periodic table providing a reference for comparing the strength of the pull on a bound pair of electrons.
- đ Lewis structures are a fundamental way to represent molecular bonding, showing valence electrons and indicating possible molecular geometries, though they may not always be the best representation.
- đ Condensed formulas provide a shorthand notation for writing compounds, omitting bonding details but showing the total number of atoms and elements.
- đ Bond line or skeleton forms are a structural representation that shows how atoms are connected, omitting some atoms for clarity and focusing on key elements of the structure.
- đ The process of creating a Lewis Dot Structure involves determining basic connectivity, counting valence electrons, adding electrons to the most electronegative atoms first, and considering multiple bonds to satisfy the octet rule.
- âïž Charged species require an adjustment in the total count of valence electrons to account for the charge, adding or subtracting electrons as necessary.
- đ Formal charge and resonance structures are advanced concepts that will be used to determine the most stable and accurate representation of a molecule's structure in future lectures.
Q & A
What is the first type of bonding discussed in the script?
-The first type of bonding discussed is ionic bonding, which involves a positively charged ion (cation) and a negatively charged ion (anion), formed by the transfer of electrons between elements.
What are some common examples of ionic compounds mentioned in the script?
-Examples of ionic compounds mentioned include potassium chloride (KCl) and sodium chloride (NaCl), as well as potassium carbonate, which involves a metal and a polyatomic anion.
What is the significance of electronegativity in covalent bonding?
-Electronegativity is significant in covalent bonding as it determines the difference in the pull of a bound pair of electrons, which can lead to polar or nonpolar covalent bonds.
What are the two main types of covalent bonds discussed in the script?
-The two main types of covalent bonds discussed are polar covalent bonds, which have a difference in electronegativity, and nonpolar covalent bonds, where electrons are shared equally due to a small or no difference in electronegativity.
How does molecular geometry contribute to the polarity of a molecule?
-Molecular geometry contributes to polarity by determining the spatial arrangement of atoms and the distribution of electron density, which can result in polar molecules if there is an uneven distribution of electron density.
What is the role of lone pairs in determining the polarity of a molecule?
-Lone pairs are areas of high electron density that can affect the overall electron distribution in a molecule, potentially leading to a dipole moment and making the molecule polar.
Can you explain the concept of symmetry in relation to nonpolar molecules?
-Symmetry in molecular structure means that the arrangement of atoms is such that any polarities are balanced and cancel each other out, resulting in a nonpolar molecule, as seen in the example of carbon dioxide (CO2).
What are the three structural representations mentioned in the script?
-The three structural representations mentioned are Lewis structures, condensed formulas, and bond line or skeleton forms, each providing different levels of detail about molecular composition and bonding.
How is the basic connectivity determined in a Lewis structure?
-The basic connectivity in a Lewis structure is determined by identifying which atoms are attached to each other, typically placing the least electronegative element in the center and connecting other atoms accordingly.
What is the octet rule, and how does it apply to the Lewis structure of the molecule discussed in the script?
-The octet rule states that atoms tend to form bonds in such a way that they have eight electrons in their valence shell, achieving a stable electron configuration. In the script, the carbon atom in the molecule COCl2 initially does not have a full octet, leading to the formation of a double bond with oxygen to satisfy the octet rule.
How do you determine the correct number of valence electrons for a molecule or ion?
-The correct number of valence electrons is determined by counting the valence electrons from each atom and adding them together. For charged species, you must also account for the charge by adding or subtracting electrons accordingly.
Outlines
đŹ Ionic and Covalent Bonding Basics
This paragraph introduces the fundamental concepts of ionic and covalent bonding in chemistry. Ionic bonding is explained as the process where a metal (cation) and a non-metal (anion) form a compound by exchanging electrons. Examples such as potassium chloride and sodium chloride are given. The importance of electron transfer in forming ionic compounds is emphasized. Covalent bonding is then divided into two types: polar and non-polar. Polar covalent bonds are characterized by a difference in electronegativity between atoms, leading to an unequal sharing of electrons. Molecular geometries and the interaction of lone pairs are highlighted as key factors in determining polarity. Non-polar covalent bonds involve equal sharing of electrons with minimal electronegativity difference, resulting in no net dipole moment. The paragraph sets the stage for a deeper exploration of these concepts in subsequent lectures.
đ Understanding Polar and Nonpolar Covalent Bonds
The second paragraph delves deeper into the nature of polar and nonpolar covalent bonds. It explains that the polarity of a bond is determined by the difference in electronegativity between the atoms involved. A visual representation using arrows to indicate electron density towards the more electronegative atom is introduced. The concept of molecular geometry and its impact on the overall polarity of a molecule is discussed, using carbon dioxide as an example of symmetrical cancellation of dipole moments. The paragraph also contrasts polar and nonpolar bonds, highlighting that nonpolar bonds occur when the electronegativity difference is minimal and the electrons are shared equally. Methane is presented as a classic example of a nonpolar molecule due to the small electronegativity difference between carbon and hydrogen. The importance of symmetry in determining molecular polarity is underscored, and the paragraph concludes with an introduction to different structural representations of molecules, such as Lewis structures, condensed formulas, and bond line or skeleton forms.
đ The Process of Creating Lewis Dot Structures
This paragraph outlines the method for creating Lewis dot structures, starting with the basic connectivity of atoms. The least electronegative element is placed in the center, and other atoms are attached accordingly. The process involves counting valence electrons from each atom and adjusting for the molecule's charge. The importance of electronegativity in determining which atoms receive electrons first is emphasized. The paragraph explains the rule of placing electrons on the most electronegative atoms and forming multiple bonds when necessary to satisfy the octet rule for atoms like carbon, nitrogen, oxygen, and fluorine. The example of Cl2CO is used to illustrate the step-by-step process of creating a Lewis structure, including the formation of a double bond to satisfy the carbon's need for a full octet. The paragraph concludes with a brief mention of resonance structures and formal charge, topics that will be revisited in future lectures.
đ Advanced Lewis Structure Construction
The final paragraph provides an in-depth look at the Lewis structure construction process, focusing on the distribution of valence electrons and the formation of multiple bonds. It explains the rule of prioritizing the most electronegative atoms when placing electrons and the subsequent need to form multiple bonds to complete the octet for certain atoms. The example of Cl2CO continues, demonstrating how to address the carbon's incomplete octet by forming a double bond with oxygen, which involves the donation of a lone pair from the oxygen atom. The paragraph also touches on the concept of resonance structures and the use of formal charge in determining the most stable Lewis structure. It concludes by summarizing the importance of understanding different types of bonds and the process of constructing Lewis structures, setting the foundation for further studies in chemistry.
Mindmap
Keywords
đĄIonic Bonding
đĄCovalent Bonding
đĄElectronegativity
đĄMolecular Geometry
đĄLone Pairs
đĄPolar Molecules
đĄNon-polar Molecules
đĄLewis Structure
đĄCondensed Formula
đĄBond Line Form
đĄOctet Rule
Highlights
Introduction to the two types of chemical bonding: ionic and covalent.
Explanation of ionic bonding involving cations and anions formed by electron transfer.
Examples of ionic compounds like potassium chloride and sodium chloride.
Discussion on metal can species and polyatomic anions in ionic compounds.
Polar covalent bonds based on electronegativity differences and molecular geometries.
Importance of electronegativity in determining the polarity of covalent bonds.
The concept of nonpolar covalent bonds where electrons are shared equally.
Illustration of polar and nonpolar bonds using molecular geometry and symmetry.
Process for determining if a molecule is polar by analyzing electronegativity differences, molecular geometry, and lone pairs.
Introduction to Lewis structures and their role in depicting molecular bonding.
Guidelines for creating Lewis structures, including basic connectivity and electron counting.
Explanation of how to account for charged species in Lewis structures.
Step-by-step process for drawing Lewis structures for a given molecule.
The significance of the octet rule and multiple bonding in Lewis structures.
How to determine which atom will donate lone pairs to form multiple bonds.
Different structural representations: Lewis structure, condensed formula, and bond line or skeleton form.
Importance of symmetry in determining the overall polarity of a molecule.
Review of the process for creating Lewis Dot Structures for a specific molecule, Cl2CO.
Final remarks on the importance of understanding different types of bonds and Lewis structures in chemistry.
Transcripts
hi students for the second video for
lecture one we're now going to continue
our general chemistry review and now
talk about the two types of bonding so
remember that we talked about our
molecular bonding in chapter 3 in your
general chemistry one course so what is
the first type of bonding that we want
to talk about well that is going to be
our ionic bonding remember that involves
a positively charged ion which we would
commonly call A
cation and a negatively charged ion so
the important part that you want to have
for the ionic compounds is that you must
form the ion this is done by gaining or
losing electrons from our elements
various examples that you would see back
in your general chemistry were such
things as potassium I chloride you would
also see such things as sodium chloride
and then you got to species that you
really didn't like and that is your
metal cans so in this case pottassium
but then your polyatomic anions an
example of that being such things as
carbonate so this being potassium
carbonate so when dealing with ionic
compounds it was normally a metal and a
non-metal one losing an electron the
other gaining it and forming a compound
based off of the charge difference this
is how ionic bond works for Cove valent
Bond we actually are going to have two
types of bonds that we need to
discuss the first of which is the polar
calent polar calent bonds are going to
be based off the difference of one of
our periodic trends that you really
needed to focus in back on your general
chemistry this is going to be the
difference in electr
negativity remember that electro
negativity is a relationship Trend in
which we are comparing the strength of
the pull of a bound pair of electrons in
a CO valent polar or non-polar Bond so
this is going to specifically talk about
the tug of war on a pair of electrons in
a bond so the question becomes how or
why does this all work well the first
thing that we need to really think about
in here is that our difference in electr
negativity
then mixed with our ideas of molecular
geometries as well as the interaction
with lone pairs
is what's going to give us our possible
polar species three very very very
important ideas that all lead to our
polar species so it's electr negativity
difference our molecular geometries then
inside of our molecular geometries the
interactions with our lone pairs finally
giving us the possibility of polar Cove
valent compounds we will talk about how
to study that in our next lecture
finally the other type of Cove valent
bond is the non-polar so what is a
nonpolar Bond well if a polar bond is
talking about the difference in electro
negativity for a non-polar Bond we are
going to share our electrons
equally this means that the difference
in electr negativity is very small and
those atoms in the bond are going to
have an equal right to those electrons
so we don't have a electron moment going
towards either one of the atoms so what
does this actually look like for us well
as we're about to see with our polar
species if I have something like
HCL we are always going to have a really
nice way of indicating which species is
going to be more polar we are going to
draw an arrow in which the arrow is
pointing towards the more electr
negative atom and a little plus sign is
going to be here above the more
electropositive atom that way you can
see that the electrons for this HCL Bond
are going to go more towards the
chlorine so in a good example of this
and talking about our molecular
geometries is if I have oxygen bound to
carbon bound to another oxygen I'm going
to form carbon
dioxide but in this case one of the
oxygens is going to pull electron
density towards that oxygen while the
other oxygen is going to pull electron
density towards itself notice that these
are going to cancel out this is
typically what we call symmetry we are
going to see a lot of these examples
when we get to our second lecture so
what does it mean for a nonpolar
compound to share electrons equally well
if I do that probably the best example
is going to be methane this is carbon
surrounded by four
hydrogens the electr negativity
difference between carbon and hydrogen
is not that large so therefore it's
going to be considered a very equal
sharing and so all of the possible
dipole moments if there are any would a
cancel out but there aren't going to be
any due to the lack of a noticeable
electr negativity difference it's really
important that we go through a very nice
seamless process in order to determine
whether something is polar or not the
best way of doing that is start with the
electr negativity difference of all of
our bonds
so this means that we are going to need
to use our periodic trend and probably
determine the calculated value in order
to see whether if an individual bond is
polar or not and then once we have done
that then we will go to the molecular
geometry I'm going to abbreviate this
just as mg molecular geometries are
going to show us that certain molecular
geometries are going to lend us to the
possibility of certain polar molecules
we will talk about those in lecture two
finally once we get done with the
molecular geometries we must go and see
if there are lone
pairs on our Central atoms lone pairs
are going to be areas of high electron
density and might lead us to having a
possible dipole negative portion of our
molecule and then finally one area that
we must go and look at is our symmetry
this was seen in the carbon dioxide
example listed up above carbon and
oxygen individually are going to be
considered polar bonds but because
they're pulling opposite and equal you
are seeing that this is going to be a
nonpolar molecule that contains bonds so
for our blue example above that's not a
polar molecule because it has dipole
moments pulling in opposite and equal
directions so now the question becomes
how can I represent all of the various
different structures that we're going to
see in this course well there are going
to be three that we will run into all
the time you know the first one this is
going to be the Lewis structure the
Lewis structure is the type of molecular
representation that we used in chem 119
the Lewis structure is going to take the
veence electrons will show very
rudimentary bonding and give us a little
bit of a possible indication when it
comes to the different molecular
geometries but this may not be the best
representation of our structure another
way of doing this is what we call a
condensed formula a condensed formula is
going to no longer show any bonding but
is going to be a very nice shorthanded
information about what atoms I have
bound to one another and going to just
show a total number of atoms and
elements in our compound this is
actually very nice when we are writing
out compounds and we were trying to put
them in paragraph form finally probably
the most important of
our structural representations is going
to be a new one for you and one that we
will practice through chapter one and
that goes by a couple of different names
probably the most common name is known
as Bond line
form but many textbooks will also call
this a skeleton form skeleton form
meaning that it's going to show you how
things are
connected and it is going to show some
very important certain atoms while going
and omitting some of the atoms for a
little bit of clarity sake so as you can
see as we go along you should be very
good with the Lew structure but I will
show you how to do that here on the next
slide while we will talk about the
condensed formula and the bondline
skeleton form when we get to lecture
two so now let's go back and review a
simple process for creating a Lewis Dot
Structure and we're going to go ahead
and do that for
cl2
Co so what we need to do in here is
first start with the basic connectivity
we need to go and figure out what atoms
are attached to each other the best way
of explaining this is that the least El
electr negativity element should be
located in the middle now what I will do
is attach an oxygen to it and I will
attach one chlorine to either side this
is the basic skeleton structure what is
so important about this is that it is
going to give us our basic
connectivity so now the question becomes
what do we do next now we need to count
the veence electrons coming from the
periodic table so when we are counting
electrons it is important to remember
that we are only using the veence
electrons because those are the ones
that are going to help us with bonding
so for our chlorines you will notice
that I have two of them they are in the
hallogen column which is going to give
us seven electrons so we will multiply
those by two chlorines and we will get a
total of 14 electrons from chlorine
alone when it comes to my carbon I am
going to get four
electrons and finally with my oxygen I
am going to get six veence
electrons and now all I'm going to do is
add this up if I take the sum of all of
my four atoms you will see that I get
24 veence
electrons this is really nice because
this is a neutrally charged species but
we must go back and review a rule that
we discussed back in our general
chemistry when we were talking
about our individual species so now with
that being the case we have to go and
look at a negatively charged species
such as NO3 1
minus what we need to do is count count
the total number of electrons I get five
from the nitrogen I get six from each of
the oxyg so if I do 6 * 3 is 18 + 5 is
23 however the correct number of
electrons for this compound is actually
24 electrons the reason why is that we
must be concerned with this negative
charge located here in is an additional
electron that we must take account for
so when we have our charged
species we must go and add or subtract
electrons due to the charges so if you
see a plus charge on your ions something
like ammonia this means that it is going
to be more positive than it expects so
you are going to
lose at least one
electron so now what will happen when I
have a negative charge well a negative
charge just as we have seen here with
our nitrate we are going to have to
add an electron so this is really
important that you go and get the
correct number of electrons so that way
when you are doing your leis dot
structures you are coming up with the
right number to help you build the
structures that you
need so now we need to go and draw in
our bonds the best rule for when adding
electrons to our skeleton structure is
that we are going to add electrons to
the most electronegative atoms first so
in this case I am going to go and redraw
my skeletal
structure and now my oxygen and
chlorines are going to get the remaining
electrons the question becomes how many
remaining electrons do I have well I
start with 24 electrons but based off of
the basic connectivity I have to go and
subtract six
electrons and now it is my job to place
18 electrons around my atoms so let us
go and do that to chlorine first the
left chlorine now has six the chlorine
here now gets six more as lone Pairs and
finally six more when it comes to the
oxygen we now have a new problem and
that is my carbon does not have a full
octet so this will now lead us into our
fourth step the fourth step is now
talking about the possibility of
multiple bonds this is going to happen
if we have
atoms with
fewer than eight
electrons this is commonly known as the
octet rule if we are dealing with
certain species especially the ones that
we are going to use in organic chemistry
you will see a lot that must obey the
octet rule carbon nitrogen oxygen
and Florine so our carbon is currently
two electrons short so what we are going
to do is go and form multiple Bonds in
order to help out that compound's
current structure and electronic
Behavior so the first thing that we can
think about is going and forming a
double bond remember that a double bond
is going to be four electrons in a
shared Bond that's not the only only one
that we can do we are also able to go to
the triple
bond that is going to be six electrons
shared
coal so now if we go back to our third
step we need to figure out which one our
chlorine or our oxygen is going to go
and donate one of its loan pairs to help
the carbon form a double bond the best
way of explaining this is the species
that have the lone pairs you are going
to donate from the one that is less
electronegative in this example it's
going to be the oxygen so I'm going to
take one of the pairs of Lone pairs of
electrons on my oxygen and I'm now going
to go in and form the multiple Bond what
my final structure is going to look like
is I'm going to have a double bond to my
oxygen and then I will have single Bond
bonds to my chlorines in which the
chlorines then have three pairs of Lone
pairs of electrons you will see in later
lectures just how we are able to
determine which structures are better we
typically call those resonance
structures and we are going to use such
rules as formal charge that we learn
back in our general chemistry
course in this video you just saw how
important it was to discuss the
different types of bonds not only ionic
bonding which uses charge different but
then the two types of calent bonds polar
and non-polar finally we went back and
Revisited an old friend from General
chemistry and we started our review of
the lewis. structure which we will use
to then build further structures in
future
classes have yourself a great day
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