Chapters 11 and 12 material on final

[Zellmer, Robert]

We had a lecturers' meeting on Friday.  I have a little more info about
what's on the final from these last two chapters.

Ch 11:

I briefly mentioned ionic liquids (Chemistry and Sustainability box in
Section 11.3).  We did not cover this in any detail and it's not on the
final.

We did not cover liquid crystals.  They are not on the final.


Ch 12:

Look at what I covered class and my notes.  Also, what I have about
solids on Carmen.  Here's a little more info.

Know how much of a particle is inside a unit cell depending on where
it sits in the unit cell.  In other words, what fraction of a particle is
within the unit cell.

Corner             1/8
Edge                1/4
Face                1/2
totally inside      1


Understand the unit cells, primitive cubic (pc, also called simple cubic,
sc), body-centered cubic (bcc), and face-centered cubic fcc).

Packing efficiency in solids (hcp, ccp/fcc, bcc, pc).

For ionic substances understand how to determine the empirical
formula from a unit cell.  Based on a description of where the ions
are placed in the unit cell or a picture of a unit cell be able to come
up with the number of cations and anions in the u.c. and thus the
empirical formula.  That's important and people often mess that
up.

For example, with NaCl we found there are 4 Na+ and 4 Cl- ions
w/in the unit cell.  That means the empirical formula is NaCl,
as it should be, and there are 4 NaCl formula units per u.c.

NaCl is said to have a face-centered cubic structure, although
it doesn't look like the fcc structure you see for a single-sized
particle (such as fcc Al, CO2, C60, etc.).  For a single-sized
particle a fcc structure has 4 particles per unit cell but has
particles only at each corner and each face.  In NaCl you can
picture it as having Cl- ions in a face-centered structure with
Na+ ions on each edge and 1 Na+ in the center.  You can
also look at it with Na+ ions in a face-centered structure with
Cl- ions on the edges and 1 in the center.  Why is it called
face-centered?  If you look at the Cl- ions they form a
fcc structure.  If you look at the Na+ ions they form a fcc
structure.  You push the fcc Na+ ions and fcc Cl- ions halfway
into each other and pull the part out where the overlap and
you get the NaCl unit cell.

For CsCl there's Cl- ions at the corners with a Cs+ in the center
of the cube.  You could also picture it as having Cs+ ions in the
corners with a Cl- in the center of the cube.  It looks like a
body-centered but it technically isn't.  It's actually a primitive
cubic.  What?  The Cl- ions form a pc structure.  The Cs+ ions
form a pc structure.  Then push a corner of the pc Cs+ ions
into the center of the pc Cl- ions and you get the CsCl structure.
In any case, with Cl- ions in each corner you get 1 Cl- ion in the
u.c. and with 1 Cs+ in the center you get 1 Cs+.  That gives an
empirical formula of CsCl with 1 CsCl fu/uc

For ZnS, the S^2- ions form a fcc structure, giving 4 S^2- ions
w/in the u.c. and there are 4 Zn^2+ ions inside the unit cell.  That
gives an empirical formula of ZnS, as it should be, and 4 ZnS fu/uc.

For CaF2, the Ca^2+ ions form a fcc structure, giving 4 Ca^2+
ions w/in the u.c. and there are 8 F- ions inside the unit cell.  That
gives an empirical formula of CaF2, as it should be, and 4 CaF2 fu/uc.

There are figures for all the above in the lecture notes and the
material you'll find on Carmen for solids (as mentioned in previous
e-mails).

There are similar things for alloys in the current textbook.  Look at
Figure 12.17 in the current edition.  In the current edition, look at
EOCEs 12.29, 12.30, 12.57, 12.58, 12.114 and Exam Prep question
12.13 for more practice in determining the empirical formula for
a solid composed of different elements.

The Chem dept has a website about crystal structures,

Visualizing Crystal Structures - OSU Chem Dept<http://undergrad-ed.chemistry.ohio-state.edu/xtal/index.html>.

Click on "Crystal Structures".  It allows you to look at the unit cells,
rotate them around, see several unit cells together, etc.  You can
also find this link at the "Notes" link on my class webpage.  Make
sure to click on "All Atoms".  Also, it can be helpful to ask for more
than 1 unit cell.  Don't ask for more than 8 uc at a time, it gets too
"busy".

You should also understand the 4 different types of solids, the
attractive forces involved and the general properties of each type of
solid (conducting or not, melting points, hardness, brittleness, etc.).
I went through this in lecture last Thursday. The following link will
bring up the summary table, which is in the lecture notes,

Types of Solids and Their Properties<https://www.asc.ohio-state.edu/zellmer.1/chem1210/notes/Table_13-10_solids_no_lines.pdf>

There aren't a lot of covalent network solids.  Pretty much memorize
the ones listed in the table.  They can be "tricky".  SiO2 is a metalloid
with nonmetal and Si is in the same group as C so you might expect
SiO2 to behave like CO2.  CO2 is a molecular solid and you might
expect SiO2 to be molecular.  That's not the case. SiO2 is a covalent
network solid.

I showed a density calculation in class.  We decided to not include
unit-cell calculations on the final exam (density, edge length, etc.).
There might be one or two density calculations in the Mastering.

Dr. Zellmer

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