Lesson 2: The Ionic Bond

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Fascinating Chemistry Lessons
Lesson 2: The Ionic Bond
Slide 1: Introduction Slide
Slide 2: The atomic number
Here are the first 20 elements of the periodic table.
The number in each box tells you how many protons
each atom contains in its nucleus. The number in the
box does not refer to the number protons and
electrons. Why not?
Because I just showed you that with the Van de Graaf
generator, removing electrons from an atom does not
change the name of the element. The only name
change when the number of electrons decreases or
increases is the atom being called an “ion” of that
element. Only if we change the number of protons in
the nucleus does the name of the element change.
What makes carbon – carbon, or hydrogen – hydrogen,
is the number of protons in the nucleus, not the
number of electrons circling the nucleus.
The number in each box is called that element’s
“atomic number.”
Slide 3: Valence electrons
The electrons in the outer ring of an atom are known as the "valence" electrons. In this picture,
the valance electrons are the ones in yellow.
The give-and-take ionic bond is typically used by atoms with only 1 or 2 valence electrons in
their outer ring looking for an atom needing 1 or 2 electrons to fill up its outer ring.
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Take an atom like sodium with 1 valence electron in its
outer ring, ring 2. Sodium has little chance of filling up its
outer ring with 7 more electrons. But when sodium spots
chlorine needing just 1 more electron to fill up its outer
ring, sodium can simply give chlorine its single electron in
ring 3. Chlorine’s outer ring will be filled, but will sodium’s?
Yes, because when sodium gives chlorine it outer electron,
sodium no longer has a ring 3. Ring 2 is now the outer ring,
and its ring 2 is filled with 8 electrons. Sodium now has a
full complement of 8 electrons in its outer ring.
Slide 4: How ionic bonds form molecules
Before sodium gives chlorine its outer electron, sodium has
an equal number of protons and electrons, and is
electrically neutral.
After sodium gives its outer valence electron to chlorine,
sodium is left with more protons than electrons, which
makes sodium electrically positive. Likewise, with that extra
electron, chlorine becomes electrically negative. So in
giving away its valence electron to chlorine, both sodium
and chlorine became ions.
What do we know about positive and negative electrical charges? They attract each other. So
when sodium gives its outer valence electron to chlorine and sodium becomes positive and
chlorine negative, the sodium and chlorine ions pull together. From here on sodium and
chlorine are bonded together as a single molecule that we call sodium chloride.
The everyday name for sodium chloride is table salt, and the bond holding the sodium and
chloride ions together is an “ionic bond.”
Sodium chloride is nothing like sodium or chlorine alone. Pure sodium will explode if placed in
water and chlorine is a poisonous gas, but bonded together, they make table salt. And that’s
true of all molecules.
You cannot predict the properties of a molecule by knowing the properties of the parent atoms.
To predict the properties of a molecule, you have to know which of the four intramolecular
bonds was used to form the molecule.
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Slide 5: Reaching a lower energy level
Sodium and chlorine bond so that they could both
lower their energy level. How much energy do the
neutral sodium atom and the neutral chlorine atom
get rid of by forming an ionic bond with each other?
To answer that question, we need to look at the
energy sodium loses when it gives up its electron and
the energy chlorine loses when it accepts sodium's
In fact, sodium gains energy when it gives up its
electron, because it takes energy to pry off sodium’s single valence electron. The energy
needed to remove sodium’s valence electron is called "ionization energy," because removing
the valence electron makes a neutral sodium atom into a positive ion.
Chlorine, on the other hand, does shed some energy when it fills up its outer ring with sodium’s
single electron. The energy given off by the neutral chlorine atom when it accepts sodium’s
electron to become a negative chlorine ion is called chlorine's “electron affinity.”
In forming sodium chloride from sodium and chlorine atoms, sodium’s ionization energy is
greater than chlorine’s electron affinity, so sodium
chloride is actually at a higher energy level than sodium
and chlorine alone.
To make it worthwhile for sodium and chlorine to
combine, sodium chloride molecules need to hook up
with other sodium chloride molecules and form a three
dimensional structure called a lattice, resembling the
monkey bars kids climb on in a playground. In doing so,
sodium chloride molecules will lose energy and reach a
calmer, more relaxed state. Here’s why.
Slide 6: Lattice energy
When sodium and chlorine bonded into a single molecule of
sodium chloride, did the sodium side of the molecule
remain positive and the chlorine side remain negative? Yes,
the sodium side remained positive and the chloride side
remained negative. A molecule that’s strongly positive on
one side and strongly negative on the other is called
“polar,” like the north and south poles of the earth, or the
north and south poles of a magnet.
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And like poles of a magnet, the positive sodium side of
a sodium chloride molecule is attracted to the chloride
side of any other sodium chloride molecule that
happens along. Polar molecules are sticky.
When two sodium chloride molecules bump into each
other, the positive sodium side of one sodium chloride
molecule sticks to the negative chloride side of the
other sodium chloride molecule.
The bond that forms between two molecules is called an intermolecular bond. “Inter” means
between, referring to between two molecules.
The bond that forms between two atoms in a molecule
is called an intramolecular bond. “Intra” means within,
referring to the atoms within a single molecule.
As more and more sodium chloride molecules happen
along, the molecules begin to stack together like bricks
forming a cube. Before you know it, the sodium and
chloride ions forget which sodium and chloride ions
they belong to. Every sodium ion is equally attracted to
all chlorine ions around it, and every chlorine ion is
equally attracted to all the sodium ions around it.
Any two sodium and chloride ions are no longer called a molecule, but a “formula unit.”
Because every sodium and chloride pair is identical, every bond holding a sodium ion to a
chloride ion becomes identical.
By settling into a rigid lattice like this, sodium chloride molecules no longer need to expend so
much energy jumping and bouncing around. The energy they no longer have to expend is called
“lattice energy.” By giving up their lattice energy, sodium chloride molecules finally reach a low,
stable energy state that more than makes up for the
energy spent removing sodium’s outer electron and
transferring it to chlorine.
The sodium and chloride ions grow into a square-shaped
crystal lattice, visible under a magnifying glass as cubic
crystals of salt.
Click the term “salt crystals” to learn how to make your
own salt crystals.
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Slide 7: Melting ionic crystals
A crystal lattice of sodium chloride is at such a low energy level and so stable that a great deal
of energy is needed to pull sodium and chloride atoms apart. If you were to heat table salt, you
would have to heat it to 1474 degrees Fahrenheit, 801 degrees Celsius to pull apart sodium and
chloride atoms.
Where, then, does the energy come from to break a salt
crystal apart when you dissolve salt in water? From the
water, which you can detect by the slight drop in water
temperature when you dissolve salt in water.
If it takes energy to dissolve salt in water, and adding
energy pushes salt’s energy state to a higher level, why
does salt ever dissolve in water?
Entropy. Changing highly organized formula units into a swirling array of sodium and chloride
ions increases the entropy. The disorder that comes when salt crystals are dissolved in water
more than makes up for the energy expended in causing it to happen. Pulling apart sodium and
chloride atoms by dissolving them or by heating them doesn’t pose a risk of releasing explosive
sodium and poisonous chlorine, because the sodium and chloride atoms separate as sodium
and chloride ions, which are not poisonous.
Slide 8: Cracking a salt crystal
And yet, you can easily crush a salt crystal with a hammer or crack
it by tapping it with a sharp knife.
Why is mechanical energy so much more effective than heat
energy at breaking the ionic bonds in a salt crystal?
When a crystal of salt is struck with a hammer, it cracks
between the rows of sodium chloride ions. Why?
The force of this hammer shoves a section of sodium
chloride formula units forward, placing positive sodium
ions right next to other positively-charged sodium ions,
and negative chloride ions right next to other negativelycharged chloride ions. The combined force of positive
sodium ions repelling other positive sodium ions, and
negative chloride ions repelling other negative chloride
ions, splits the crystal. The split continues down the plane between two rows of sodium
chloride formula units, cracking the crystal and causing it to shatter. The term for cracking
easily when subjected to mechanical force is “brittle.”
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Slide 9: Catching our breath
Let’s stop for a second and look back.
Ionic bonds typically form between atoms looking to give
away 1 or 2 valence electrons and atoms looking for 1 or
2 electrons to fill up their outer ring.
When one atom does give away its valence electron to
another atom, both atoms become oppositely-charged
Oppositely-charged ions attract one another and form an ionic
Ionic bond molecules are polar. By attaching all their opposite
polar ends together, ionic molecules form crystal lattices,
a process that sheds them of lattice energy and result is a
heat-stable crystal. Each individual molecule in a crystal
lattice, being equally attracted in all directions to other
identical molecules, is no longer called a molecule, but
rather a “formula unit.” Crystal lattices resist heat, but
crack easily when struck, a property known as “brittle.”
Slide 10: Ionic bonding between rows of the
periodic table
Sodium was able to bond ionically to chlorine. Do give-and-take ionic bonds require the atoms
to be in the same row of the periodic table?
Could lithium in row 2 have given its electron to chlorine in row 3 needing one electron to fill its
ring 3?
Of course. Lithium could give its single electron in ring 2 to chlorine’s open slot in ring 3. Lithium
would become a positive ion
and chlorine a negative ion
and the opposite electrical
charges would draw them
together into a give-and-take
ionic bond just like lithium
and fluorine.
The molecule is called lithium
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Slide 11: Ionic bonds involving two electrons
Give-and-take ionic bonds can also form between an
atom with 2 electrons in its outer shell and an atom
needing 2 electrons to fill its outer shell. In row 2 of the
periodic table, for example, beryllium has 2 electrons in
its outer ring, ring 2.
Oxygen needs 2 electrons to fill its ring 2. When
beryllium gives its two electrons to oxygen to form
beryllium oxide, beryllium becomes a 2 plus positive ion
and oxygen a 2 minus negative ion. This difference in
electrical charge pulls the ions together into a single polar
Beryllium oxide is a hard ceramic material
used in spark plugs because it is so resistant
to heat.
Slide 12: Giving electrons without ionic bonding
Ionic bonds form when one atom gives its electrons and the other atom takes the electrons, but
just because one atom transferred its electrons to another atom does not mean that an ionic
bond has formed. The ionic nature of a bond refers to how the two atoms share each other’s
electrons once they have bonded together. If one atom does completely take control of the
other atom’s electrons, the bond is ionic.
Is there any way to predict how the two atoms will
share their electrons inside a bond? Aluminum, for
example, forms an ionic bond with nitrogen after giving
nitrogen its outer three electrons.
Boron, right above aluminum in the periodic table, also
appears to give nitrogen its three outer electrons, but
boron does not form an ionic bond with nitrogen
because boron retains some control of the electrons it
is offering to share with nitrogen.
Could this have been predicted?
Slide 13: Pauling’s electronegativity chart
In 1932, Linus Pauling offered a chart to predict how atoms would share their electrons. The
chart lists the “electronegativity” of each element, indicating its attractive power for electrons
in a chemical bond.
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Since the basis for electron attraction begins with the
number of protons in the nucleus, in any row of the
periodic table, electronegativity increases as you add
more protons to the nucleus.
However, electronegativity decreases in every column of
the periodic table as the outer electron moves further
and further from the nucleus, and more and more
electrons accumulate to block the nucleus’ pull on the
outer electrons.
Electronegativity resembles ionization energy and
electron affinity, but those two terms refer to removing
and adding electrons to single isolated atoms before
they’ve bonded. Electronegativity refers to the tug of
war between atoms after they’ve bonded.
Pauling said that the polarity of an intramolecular bond
could be predicted by the electronegativity difference
between the two atoms forming that bond. An
electronegativity difference of greater than 2.0, he said,
would produce a give-and-take ionic bond.
Subsequent chemists have modified this a bit for metals.
For example, in our example of aluminum nitride, the
electronegativity difference is 1.43, 3.04 for nitrogen
minus 1.61 for aluminum, but because metals, like
aluminum, have a looser hold on their electrons than
non-metals, aluminum nitride is ionic even though the
electronegativity difference is less than 2.0.
Likewise, when three magnesium atoms bond to two
nitrogen atoms, the electronegativity difference for each
bond is 3.04 minus 1.31, or 1.73, suggesting the bond won’t be ionic, but magnesium, being a
metal, is quite willing to release its outer electrons and does form an ionic bond with nitrogen.
Slide 14: Alchemy
When sodium gave away its outer electron to chlorine, did
sodium and chlorine become different elements now that
they have different numbers of electrons?
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No. The only thing that identifies an element is the number
of protons in its nucleus. So long as that does not change,
the element does not change. The only thing that does
change is that the element changes from a neutral atom to
an ion of that element.
So when a van der Graaf generator strips an electron off
each of these atoms, the atoms simply become positive ions
of that element, because the number of protons in the
nucleus hasn’t changed.
For centuries, alchemists tried to turn lead into gold. They
were never successful, because there was no way to remove 3
protons from lead’s nucleus and decrease its atomic number
from 82 to 79, the number of protons in gold’s nucleus.
Slide 15: Atomic Size
Here are the first 20 elements in the periodic table with the number of protons for each
element. Each row, or period, indicates a new ring of electrons. Hydrogen and helium have only
1 ring. Lithium, beryllium, boron, carbon, nitrogen, oxygen, fluorine, and neon have two rings.
Elements sodium through argon have three rings, and
potassium and calcium have four rings.
Each column of the periodic table tells us how many
electrons are in the outermost ring. So, for example,
lithium in the first column has 1 electron in its ring 2.
Beryllium in the second column has 2 electrons, boron in
column 3 has 3 electrons, and so on.
As you can see, adding a proton increases the size of the
Does adding electrons increase the diameter of the whole atom?
Let’s see. Which atom do you think is bigger: hydrogen or helium?
Hydrogen. The diameter of a hydrogen atom is 74 picometers. A picometer is 1 trillionth of a
meter. Hydrogen is 74 picometers wide, but an atom of helium is only 64 picometers wide. Why
should that be?
Why should helium with 2 electrons be smaller than hydrogen with only 1 electron? There must
be something else besides the number of electrons that determines the size of an atom.
There is. The number of protons in the nucleus. With each additional proton, the nucleus is able
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to attract the electrons around it with greater and greater force. As the positive electrical force
in the nucleus increases, the nucleus is able to pull the electrons in each ring closer to the
nucleus, making the atom with more electrons smaller than the atom to its left.
Let’s see how this plays out in the first two rows of the periodic table. The number above each
atom indicates the diameter of the atom. Notice how, in each row, the atoms tend to get
smaller as electrons are added to the outer ring. As the nuclei gain protons, the nuclei are able
to pull the electrons inward with more and more force.
Why, though, is the difference in size between
lithium and beryllium, in row 2, so much more than
the difference in size between carbon and nitrogen?
Lithium is 304 picometers in diameter and beryllium
is only 222 picometers, a difference of 82
picometers. Carbon is 154 picometers in diameter,
but its neighbor nitrogen is only 14 picometers
smaller at a 140 picometers in diameter.
Remember that electrons repel each other. Even if a larger nucleus is able to pull its electrons
toward it with more force, as the electrons try to move toward the nucleus, they also begin to
repel each other. The more electrons, the more they repel each other. So the electrons in
carbon and nitrogen, with 4 and 5 electrons in ring 2, can’t pull together as much as beryllium
with only 2 electrons in ring 2. In fact, look at oxygen and fluorine. Oxygen’s electrons and
fluorine’s electrons repel each other so much, that oxygen and fluorine are actually larger than
Slide 16: Electron Affinity
Electron repulsion has a significant effect on both ionization
energy and electron affinity. Ionization energy, you recall, is
the energy needed to remove a valence electron from a
single atom. Electron affinity is the opposite. Electron affinity
is the energy needed to add a valence electron to a single
Ionization energy is always a positive number, because it always takes energy to remove an
electron from a single atom. Electron affinity, however, can be positive or negative, because
many times adding an electron releases energy and makes the atom more stable. When energy
is released by adding an electron, the electron affinity is a negative number.
Let’s see how electron repulsion affects ionization energy and electron affinity. Each row of the
periodic table is called a “period,” while each column of the periodic table is called a “group.”
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In the first column of the periodic table, group 1, lithium
is situated in the second row, period 2. It doesn’t take
much energy to remove lithium’s single electron in ring
2, because lithium wants the electron removed since
lithium could then exist with a filled outer ring, ring 1.
Notice that it takes even less energy to remove the
single electron from sodium’s ring 3. The reason for this
is that the 8 electrons in ring 2 are shielding the valence
electron in ring 3 from the nucleus’ positive electrical
charge and even helping to repel the valence electron
away from the nucleus.
The same goes for potassium, but its single valence electron in ring 4 is even further from the
nucleus with even more electrons shielding the valence electron from the nucleus’ strong
electrical attraction while simultaneously repelling the electron away from the nucleus.
Over to the right in group 7, it’s understandable why fluorine has such a high ionization energy.
Fluorine is looking for an electron to fill up its ring 2. The last thing it wants is to lose one of its
outer electrons.
Chlorine has less of a grip on its outer electrons because its outer electrons in ring 3 are being
shielded by the 8 electrons in ring 2. Bromine has an even weaker attraction for its outer
electrons because of all the electrons between the nucleus and ring 4. When we flip this
around and look at electron affinities, similar logic applies.
In the first column of the periodic table, group 1, the positive electron affinity for lithium means
that it takes energy to add an electron to lithium. That’s because lithium is anxious to rid itself
of its single electron so it can exist with its outer ring filled.
Why, though, is the electron affinity for sodium lower
than for lithium? Why is it easier to slip an extra
electron in to sodium’s outer ring than into lithium’s
outer ring? Because the eight electrons in ring 2 are
shielding sodium’s nucleus from the effect of another
electron around it.
Likewise, potassium’s nucleus is even more protected
from the addition of one more electron, making it
even easier to slip an extra electron into potassium’s
ring 3.
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The interesting thing about electron affinity occurs with fluorine. The negative value indicates
that fluorine sheds energy when it finds an 8th electron for its ring 2, which is understandable as
that 8th electron fills up fluorine’s outer ring.
What seems out of kilter, though, is that chlorine sheds even more energy than fluorine when
an 8th electron slips into its outer ring. You’d think that with 8 electrons in chlorine’s ring 2,
those eight electrons would be shielding the ring 3 electrons from chlorine’s nucleus and would
even be repelling chlorine’s electrons in ring 3 away from the nucleus. You would expect the
eight electrons in chlorine’s ring 2 to lower chlorine’s electron affinity, not raise it. What gives?
Electron repulsion – in fluorine. The 8th electron being added to fluorine’s ring 2 has to squeeze
into a tiny orbit because, as you recall from the last few, fluorine is a very small atom. Trying to
squeeze into such a small orbit brings fluorine’s 8th electron into close vicinity with the other
seven electrons in ring 2. The eight electrons end up repelling each other and making it more
difficult for that 8th electron to join fluorine’s ring 2. The result is that fluorine’s electron affinity
is actually less than chlorine’s.
Slide 17: What you know so far
1. The atomic number indicates the number of protons in the element’s nucleus and the
number of electrons orbiting around the nucleus.
2. Elements are defined by the number of protons in their nucleus (the atomic number), not the
number of electrons around the nucleus.
3. When there are more electrons than protons, the atom is a negative ion; fewer electrons
than protons, a positive ion.
4. The electrons in the outer ring of an atom are the valence electrons.
5. Atoms bond to other atoms to shed themselves of excess energy.
6. Atoms bond to other atoms by manipulating their outer valence electrons in such a way that
they end up with their outer rings filled with electrons, even if it means emptying the outer
Slide 18: What you know so far
7. Atoms bond ionically when one atom gives one or more outer valence electrons to another
atom. The giving atom becomes a positive ion in the process, while the taking atom becomes a
negative ion, and since opposite charges attract, the two atoms pull together.
8. The electrons being given do not have to end up on the same ring in the taking molecule as
they started out from on the giving molecule.
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9. The properties of a molecule rarely if ever resemble the properties of the atoms making up
the molecule.
10. While it takes energy to remove electrons from the atom giving away its electrons, energy is
shed by the taking atom and still more energy, called lattice energy, is shed when the now polar
molecules form a crystal lattice of formula units.
11. Ionic molecules in a crystal lattice are no longer called molecules, but “formula units.”
Slide 19: What you know so far
12. The bond holding two atoms together in a molecule is the intramolecular bond, while the
bond attracting one molecule to another is the intermolecular bond.
13. Crystal lattices crack readily when one row of the lattice is shoved forward and is now
aligned with its ions next to other like-charged ions, causing the two rows to repel apart.
14. Crystal lattices are resistant to heat, meaning they have a high melting point.
15. Pauling’s electronegativity chart quantifies the attraction each atom in a molecule has for
electrons. When the difference between the two atoms reaches 2.0, the intramolecular bond
holding them together is likely to be ionic.
Slide 20: What you know so far
16. In each row (period) of the periodic table, the atomic diameter decreases as more protons
are added, because the electrons are pulled inward with more and more force. Each additional
electron that’s added, however, repels the other electrons and resists the nucleus’ inward pull.
17. Ionization energy is the energy needed to remove a valence electron from a neutral atom.
18. Ionization energy is low for atoms with a single valence electron (elements in group I of the
periodic table), and drops further in the larger group I elements as the valence electron is
located further and further from the nucleus.
19. Electron affinity is the energy needed to add a valence electron to an atom, or the energy
released when the valence electron is added, in which case electron affinity is negative.
20. Electron affinity has a high negative value in atoms in the group VII elements with seven
valence electrons and needing only one electron to fill their outer ring with eight electrons.
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Chemistry Lesson 2 – Lab
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