6.02 Covalent bonds

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Del curso dictado por University of Kentucky

Chemistry

498 calificaciones

University of Kentucky

Chemistry

498 calificaciones

This course is designed to cover subjects in advanced high school chemistry courses, correlating to the standard topics as established by the American Chemical Society. This course is a precursor to the Advanced Chemistry Coursera course. Areas that are covered include atomic structure, periodic trends, compounds, reactions and stoichiometry, bonding, and thermochemistry.

De la lección

Week 6

This unit looks more in depth at molecular compounds to see how they are bonded together by looking at Lewis structures. These structures provide information about the types of bonds (single, double, or triple) as well as the connectivity of atoms. By knowing the Lewis structure, we can also predict the three-dimensional geometry of an individual molecule.

6.01 Lewis symbols 4:36

6.02 Covalent bonds 5:19

6.03 Electronegativity 11:21

6.04 Lewis Structures Part 1 13:11

6.04 Lewis Structures Part 2 18:49

6.04a Draw the Lewis structure for H2O 1:13

6.04b Draw the Lewis structure for H2CO 2:00

6.04c Draw the Lewis structure for CCL4 1:21

6.04d Draw the Lewis structure for NH3 1:16

6.05 Resonance & Formal Charge 16:27

6.05a Determine the formal charg on atoms in NH4+ 2:24

6.05b Determine the formal charge on atoms in H2CO 3:17

6.06 Exceptions to the Octet Rule 8:48

6.06a Draw the Lewis structure for I3- 1:16

6.06b Draw the Lewis structure for ClF4 1:41

6.06c Draw the Lewis structure for XeF4 1:23

6.07 VSEPR Part 1 11:01

6.07 VSEPR Part 217:07

6.07a Determine the electron pair and molecular geometries for I3- 1:42

6.07b Determine the electron pair and molecular geometries for NH3 1:22

6.07c Determine the electron pair and molecular geometries for ICl4- 1:57

6.07d Determine the electron pair and molecular geometries for PF5 1:52

6.07e Determine the electron pair and molecular geometries for XeF2 2:16

Conoce a los instructores

Dr. Allison Soult
Lecturer
Chemistry
Dr. Kim Woodrum
Sr. Lecturer
Chemistry

In this module, we're going to look at covalent bond formation.

By the end of this module, you should be able to describe how covalent bonds form.

When we look at covalent bonding,

what we see is that we have electrons being shared.

And the key part about this is,

is that they are not being transferred from one atom to the other and

both electrons basically claim ownership of those electrons in that bond.

Because they’re sharing electrons, they’re not always shared evenly.

And so when we have an ionic compound where we get positive and

negative charges because somebody gained an electron, somebody lost an electron.

When we look at covalent bonding what we see is that we have a partial positive or

a partial negative charge on a molecule because of uneven sharing of electrons.

And we'll talk more about that and

how we predict that when we look at electronegativity.

The key thing that we have to worry about for

covalent bonding is that it's sharing of electrons.

Covalent compounds form between two or more nonmetals.

We also know these as molecular compounds.

The reason for this is because if I look at covalent or molecular compounds.

They exist as discrete molecules rather than an array as do ionic compounds.

For example, if I look at carbon dioxide or CO2, if I were to

zoom in on the sample, what I would see are many, many atoms of carbon dioxide.

If I did the same thing and zoomed in on sample of sodium chloride,

what I would see is an array where I had alternating sodium and chloride ions.

And while there would be a one to one ratio,

it doesn't mean there's only one sodium ion and only one chloride ion.

When I look at CO2 however, I see that I have one carbon and

two oxygens in each molecule.

We also see covalent bonds forming with elements that are the same and this helps

explain why we have some elements that exist as diatomics in nature.

We see that we can have single bonds, double bonds and triple bonds.

In a later module we'll be able to figure out why we form singovers as double versus

triple bonds, as well as, how the atoms are connected to one another.

For example, why is carbon in the middle and not an oxygen atom in the middle.

So when we have a covalent bond, it's the sharing of electrons between two atoms.

So our simplest covalent and molecular compound will be H2.

Hydrogen has one electron on each atom and when it comes together,

it now shares those two electrons between the two atoms of hydrogen.

And instead of representing this as a two dots to represent the pair of electrons,

we can now represent this as a line, which indicates a single bond.

And it's two electrons being shared.

When we get in other molecules we start to see that we have

some electrons involved in bonds, and here we have double bonds.

And since each line or each bond represents two electrons,

each double bond contains four electrons.

So that means eight electrons in carbon dioxide are involved in bonding, but

we also see that we have other electrons that are not involved in bonding, and

these are known as lone pairs, or nonbonding pairs.

They are still part of the molecule.

They're still a part of the whole structure, but

they are not involved in the bonding in this case between carbon and oxygen.

And so we call these lone or nonbonding pairs.

And we also see that we have eight nonbonding electrons or four lone pairs.

Now in this case, it's just a coincidence that the number of

nonbonding electrons and the number of bonding electrons is exactly the same.

It's not always going to be that way.

It's just the fluke of this particular example.

So we looked at Lewis symbols for individual atoms.

When we have a compound, what we look at are something called Lewis structures, and

what they do is they help us understand how atoms are connected to one another,

and what type of bonds are formed between those atoms.

Now, Lewis structures are fairly simple model,

we have a very simple set of guidelines to follow in creating those Lewis structures.

And this model does a pretty good job of explaining what's actually seen in nature.

Remember that Lewis structures are just a way of describing what

we've already seen in nature.

And we, so we see the bond type, but we also see the lone pairs of electrons.

And where this is going to become particularly important is when we start to

look at the geometry of molecules.

While the lone pairs don't effect anything about the number of bonds that we have,

those low pairs do effect the geometry or shape of the molecule.

And so we have to remember that they're there, so when we start looking at

the 3D structure molecules, we can predict that was a high level of accuracy.

In the next module, we're going to look at electronegativity which helps us

predict how electrons are going to be shared between two atoms

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