This is an introductory course for students with limited background in chemistry; basic concepts involved in chemical reactions, stoichiometry, the periodic table, periodic trends, nomenclature, and chemical problem solving will be emphasized with the goal of preparing students for further study in chemistry as needed for many science, health, and policy professions..
Ionic Bonds
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Introduction to Chemistry: Reactions and Ratios
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从本节课中
Matter and Energy
<p>If you are interested in significant figures in more detail, here are some <a href="https://www.khanacademy.org/math/pre-algebra/decimals-pre-alg/sig-figs-pre-alg/v/significant-figures" target="_blank">good videos</a> to follow on Khan Academy.</p><p>This week we will continue our explorations of matter and energy. We will discuss the sub-atomic particles that govern chemical reactions, isotopes, anions, and cations. We will learn how to name compounds, calculate formula masses, convert between grams and moles, examine periodic trends, and more! An advanced problems set is posted now; that is a longer assignment and is optional unless you would like to be eligible for the Honor’s Track. You can still earn a regular verified certificate without completing the advanced problem sets, so please be sure to keep working on the normal weekly exercises.</p>
与讲师见面
Prof. Dorian A. CanelasAssistant Professor of the Practice
Chemistry
This lecture will be about monotonic ions
and the energetics of ionic bond formation.
Between the last lecture and this lecture, I hope
you've watched the demonstration on alkaline metals reacting with water.
Besides being a fun experiment, that particular demonstration is
very illustrative of cation formations for metals that react
with water. That demonstration also introduces the
idea of periodic reactivity trends for metals in a group.
It also showed the trend, as one traverses from left to right across a period.
In the last lecture I defined the term ion, and we did some calculations
about the number of electrons and protons in different types of monatomic ions.
In this lecture, we're going to learn how to use proper nomenclature to
name ions and we're also going to talk about ionic bond formations.
Let's begin by going through
the rules for naming monatomic ions. It's exactly what it
sounds like. Monatomic means there's one atom that
composes the ion. If we're naming something that gains
electrons to become an anion, the way that it is named is simply
to change the ending of the name of the atom to -ide.
I think the easiest way to show this is by giving an example.
In the case of sulfur, it would like to gain two electrons.
It goes from being a sulfur atom to being a sulfide ion.
So the ending of the name, -ur, is replaced by -ide.
It's even simpler to name cations. When a substance that is a
main group element loses electrons to become a cation, the
name of the monatomic cation has no change from the name of the atom.
We just add the word ion after the name of the atom.
For example, in that demonstration, you saw a
sodium atom losing an electron to become a sodium ion.
Cations don't have a special ending to the name of the atom.
You could say either ion or cation, both would be correct.
Now it's your turn to practice.
If you have a capability to do so, go ahead and get a piece of paper out.
Write down these four examples. Al 3 plus, S R
two plus, N three minus, and I one minus.
And attempt to name those using the rules we've just covered.
Here's one for you to go ahead a try online.
Wonderful. Naming the cations was really easy.
The aluminum 3 plus is just called an aluminum ion,
the strontium 2 plus is just called the strontium ion.
Please note that for the main group elements, we don't have
to say anything about the charge when we name the ion.
Aluminum almost always loses all three of its electrons when it forms an ion.
Strontium almo, almost always loses two electrons when it forms an ion.
That's because these atoms and many other atoms
would like to be isoelectronic with the noble gases.
So they lose the number of electrons that will make them have the same number of
electrons as the closest noble gas or inert gas on the periodic table.
If you don't have a periodic table out, it's
always good to have one out during these video lectures.
So you might want to go ahead and get that out and look at aluminum.
And see how many electrons it has when it loses three electrons total.
The anions are a little bit more difficult to name, because we have
to remember to name, to change, I'm sorry, the ending of the name.
So N 3
minus becomes the nitride ion, and I 1 minus becomes the iodide ion.
It's always a good idea in science to ask yourself why
something is happening rather than just trying to memorize what happens.
In the formation of ions, the reason why something is happening is that there's a
driving force to form ions The electrons want to lower their energy.
Remember,
systems in nature spontaneously go to lower free energy whenever they can.
And in this case, an electron can either leave
one type of atom or join another type of atom to lower it's potential energy.
That is the driving force for ionic bond formation.
That's why ionic bonds form.
One way an electron can lower it's energy, is by getting closer to the nucleus.
Let's consider some simple drawings of hydrogen atoms.
So I'm just going to go ahead and try using some different colors here.
Remember a hydrogen atom has one proton as a nucleus and then there's one electron
that is in the space around that nucleus. I could draw
a second hydrogen atom. I'll try to use a different color here.
Still got one proton in the nucleus.
This time I'm going to draw the electron a little bit farther away.
So I have go two hydrogen atoms here.
We'll call this hydrogen b, and we'll call the other one hydrogen a.
Which of those two hydrogens has an electron with lower energy?
To give you a hint, think about what you know about opposite charges.
Do they attract each other, or do they repel each other?
And think back to those graphs of
opposite charges approaching one another.
Via Coulombic attraction and determine which of these do you think is lower
energy, the blue one, which is hydrogen a, or the red one, which is hydrogen b.
Great!
You've probably determined that the electron in
hydrogen a has lower energy than the electron
in hydrogen b, because the electrons like to
be close to protons because they're oppositely charged.
Now there is a shorthand way for drawing electrons versus energy,
so we can imagine here that energy is, free energy is
increasing as I get further away from the nucleus in these atoms.
Right?
So this is higher energy if I'm way out here and this is lower energy.
At this point, that's higher energy. This one is lower energy.
As you can see it takes up a lot of space
to draw even a simple hydrogen atom iIf I draw the
entire orbital where the electron can occupy.
So there is a shortcut that chemists use.
Chemists like to take a little slice of the orbital, like this, that's white.
And you probably can't see it.
So I'm going to go back to using green. Chemists like to take a slice like this.
So remember energy is high
at the top and low close to the nucleus.
And then instead of showing the electron as an e
minus the electron will be shown as a little arrow.
Now so if I went over here and said representations of electrons right?
Sometimes electrons are written out or you could have e minus.
Sometimes people draw an arrow for an electron.
It can be a single headed arrow or a double headed arrow.
It could
be right side down, right side up, or upside down.
All of those are ways of drawing an electron.
If I just want to show the slice that I've drawn there.
I could draw energy going from low to high, and remember
this is free energy it's not kinetic energy, or something like that.
And now I'm going to show instead of drawing that
gigantic circle for where the electron is I'm just going to draw
a little line, see how I got that, that slice there?
So this is just a line really the orbital continues around this way.
But I'm just showing a line to represent that in the
nucleus is implied its down here, but usually it's not drawn.
And to show the electron I'm just going to draw an arrow.
It can be a single headed or a double headed arrow.
So this type of drawing is used frequently where all that is shown in the drawing is
the axis for the energy, a line, and the electron.
And we can show relative electron energies using this type of picture.
Let me show you an example of what I'm doing here.
So here's a drawing of two different types of atoms.
I've shown that there are different types of atoms here with different colors.
In both cases, there is a nuclei down at
the bottom, so a nucleus with a bunch of protons.
Let me make that one red.
So there's a bunch of protons here.
This isn't necessarily hydrogen, and over here there's a bunch of protons.
That's usually left off the picture.
Can you see that? Now, free energy
is on the y axis, so high energy is up here, and low energy is down here.
And this electron is farther away from its
nucleus than the blue electrons are, you see that?
Okay, so the red electron here has higher energy.
It turns out that metals have higher energy valence electrons than nonmetals.
So the metal valence electron might be up at the high energy, as I've shown here.
And the nonmetal valence orbitals would be down
at lower energy as I've shown here in blue.
The atom I'm representing here for the nonmetal is fluorine.
And the atom that I'm representing for the metal is sodium.
So this is a picture of sodium and fluorine before they bump into each other.
They're about to have an encounter.
The sodium is about to bump into the fluorine.
And when they have the encounter, let me go back a slide.
When they have the encounter the sodium electron, oops,
these two atoms are about to have an encounter.
And when the metal bumps into the nonmetal the electron of the metal
can go to lower energy if it least the metal and joins the nonmetal.
So that's exactly what happens. When that happens
the sodium atom becomes a sodium cation, and the
fluorine atom becomes a fluoride anion. Do you see what happened?
Remember the electrons prefer to be at the lowest possible energy.
So that is the driving force from the electron that used to
be with the sodium, leaving to join the fluorine to make fluoride.
The compound that's formed is called sodium fluoride.
It's an ionic compound, because it is comprised of two
ions; a cation, the sodium, and an anion, the fluoride.
When an ionic bond forms, one atom is taking the electrons from the other atom.
They're not sharing the electrons anymore.
In fact, from the sodium's perspective, it's just lost its electron.
It doesn't necessarily know where the electron went, it's just gone.
So in an ionic bond, the electrons are taken, they're not shared, just like James
Bond's Martini is shaken, not stirred. I know, it's a bad pun, but I enjoy it.
So to summarize what just happened, when there is an encounter between a metal
and a nonmetal, let's think about what's happening to the metal and the nonmetal.
The metal's valence electrons are relatively high in potential energy.
They want to go to lower potential energy.
That's the driving force of their reaction.
The nonmetal on the other hand, unless it's a
noble gas, in which case it probably won't react.
But let's suppose it's a different type of nonmetal, not a noble gas.
It has some vacancies in its orbitals
at relatively low potential energy.
So it has some space for another electron to join the atom.
So what happened when the neutral
sodium atom encountered the neutral fluorine atom?
Can you type in your own words what happened
when the neutral sodium atom encountered the neutral fluorine atom?
Thank you for making an effort to describe
what happened during that encounter between those atoms.
Many of you said that the encounter resulted in an
electron transfer, and that that electron transfer resulted in ions.
The electron left the metal to leave a cation,
and it joined the nonmetal to make an anion.
Those two ions then stick together because of Coulomb's law.
Coulombic's attraction. And that forms an ionic compound.
Remember the Coulombic attraction is an electrostatic force, and
an ionic bond is in fact, just an electrostatic force.
One of the atoms has taken the electrons from the other atom.
And then the positive charge and the negative charge
just like to stick together because of Coulombic attractions.
That's what makes an ionic compound.
Many ionic compounds form three dimensional crystal lattices.
And many of them are also solids at room temperature.
Here's an example.
Table salt, which is sodium chloride.
You're probably very familiar with table salt in your everyday life.
In a piece of table salt, in a crystal of table salt, there are actually no discrete
sodium chloride molecules.
In other words, the bond is not directional.
The sodiums here are shown as orange.
They've lost their electrons and they tend to be slightly smaller than the chlorides.
All of these have a positive charge of course.
So, I should add the positive charge and
those are surrounded by the negatively charged chlorides.
But it's impossible
for me to say, which chloride is bonded to this sodium right?
It's surrounded actually by six chlori, chlorides.
So, ionic bonds are non-directional.
Ionic compounds, in fact, are what's called an extended solid.
We have a sodium bonded to a chlorine on once, a chloride excuse me on one side.
And on the other side, so there's
not really a directional bond.
And then this is showing just a little piece of the crystal.
There's actually another sodium right here that's
not shown, and then there's another chloride.
Do you see how that works? Wonderful.
I hope you've enjoyed learning about ion formation and ionic bonding today.
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