Transcript Periodic Table Trends

In This Lesson:
Periodic Table
Trends
(Lesson 3 of 4)
Stuff You Need:
Periodic Table
Today is Tuesday,
February 28th, 2017
Pre-Class:
Get your periodic table. Tell me some patterns
you see. They can be obvious…
Places to Look: Atomic number, atomic mass
Things to Guess: Atomic “size,” nuclear charge,
location of metals, gases, liquids?
Today’s Agenda
• Day 1:
–
–
–
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Introduction to the Periodic Table
Why does periodicity work?
Coloring the Periodic Chart
Meet the “Families”
• Day 2/3:
–
–
–
–
Chemical Periodicity
Periodic Table Trends
Alien Periodic Table
Periodic Law
• Where is this in my book?
– P. 160 and following…
By the end of this lesson…
• You should be able to use the periodic table as
a “map” to make predictions and inferences
about various elements.
• You should be able to describe valence
electrons and specify any representative
element’s number of valence electrons.
Mendeleev’s Periodic
Table
Dmitri Mendeleev
Modern Russian Periodic Table
Mendeleev and the Table
• Mendeleev organized his
table at a time in which
not all of the elements
were discovered or some
properties were
unknown.
• It was his discovery of
the patterns in the table
that allowed him to
make predictions about
the elements that would
later will the gaps in his
table.
Periodic Law
• Out of Mendeleev’s discoveries and
organization came the concept of the Periodic
Law.
• “Many physical and chemical properties of
elements recur in a systematic manner with
increasing atomic number.”
– Plain English on next slide…
Periodic Law
• In other words, there is a pattern of physical
and chemical properties to the periodic table,
and it repeats every so often.
• In fact, it repeats every 2, 8, 18, 32 elements.
– Example: Hydrogen (A# 1) is similar to Lithium (A#
3) is similar to Sodium (A# 11) is similar to
Potassium (A# 19) is similar to Rubidium (A# 37) is
similar to Cesium (A# 55) is similar to Francium
(A# 87).
– 1-3-11-19-37-55-87
Periodic Table Set Up
• Rows on the table are called
periods.
• Columns on the table are
called groups or sometimes
families.
– Elements in the same group
behave chemically similarly.
– Elements in the same period
may not behave chemically
similarly.
G
R
F O
A U
M P
I S
PERIODS
L
I
E
S
Group
Groups and Periods
Period
Valence Electrons
• We’ll do more on this one soon, but for now
we should be aware of the very basics of
valence electrons.
• Valence electrons are only the electrons in the
highest energy levels (n levels).
• Electrons in d and f sublevels don’t count.
– Electrons are the particles that matter in bonding
and really give an element its behavior.
• You can diagram out the configuration.
Valence Electrons
• Aluminum:
• 1s2 2s2 2p6 3s2 3p1
– 3 valence electrons.
• Sodium:
• 1s2 2s2 2p6 3s1
– 1 valence electron.
• Alternatively, you can use this handy pattern:
2 Valence Electrons
1 Valence
2 Valence
Electron
Electrons
5 Valence
6Electrons
Valence
8Electrons
Valence
Electrons
Electrons
3 Valence
4 Valence
7Electrons
Valence
Electrons
Roman Numerals
• Now’s probably a good time to let you know
that many periodic tables list group numbers
in Roman numerals.
•
•
•
•
I=1
IV = 4
V=5
VI = 6
• You’ll need to know up to VIII (8).
• There’s also X (10), L (50), C (100), D (500),
and M (1000).
REPRESENTATIVE ELEMENTS
Periodic Table Info
• The periodic table is arranged in its funny
shape for many, many reasons.
• For a set of elements that can at times be so
different, they follow some amazingly
predictable patterns.
• The most general pattern is the organization
between metals (1), nonmetals (2), and
metalloids (3).
An Outline Look (color this)
http://web.buddyproject.org/web017/web017/metals.html
1. Metals
• Metals are:
– Good conductors of
heat/electricity.
– Malleable (able to be
hammered – opposite of
brittle).
– Ductile (able to be pulled
into wires).
– High tensile strength.
– Shiny (have high luster).
Metal Trends
• Metals get more metallic as you go down the
groups.
• Metals get less metallic as you go across the
periods.
• In other words, in Group IA, we’d expect
Potassium (K) to be a better conductor, shinier,
more malleable, et cetera, than Lithium (Li).
• That’s why gold is generally shinier than silver.
ALKALI METALS
Alkali Metals
• Alkali metals:
– Have one valence electron.
– Are never found pure in
nature (too reactive – see
video).
– Become less reactive down
the group.
Potassium (K)
reacts with
water and must
be stored in
kerosene
Aside: Francium (Fr)
• I can’t show you a picture of Francium for this
slide.
• The reason? It’s really really rare.
• According to Bill Bryson in A Short History of
Nearly Everthing, there may be as few as 20
naturally-occurring atoms of Francium on the
Earth at any given time.
– Scientists can only guess what a large sample
would look like.
ALKALINE EARTH METALS
Alkaline Earth Metals
• Alkaline earth metals:
– Have two valence electrons.
– Are less reactive than alkali metals (but are
still not found pure in nature).
– Tend to be basic (alkaline).
Aside: Strontium (Sr)
• Between 1959 and 1970, scientists in St. Louis launched the
Baby Tooth Survey in which children’s primary (baby) teeth
were chemically analyzed.
– Over 300,000 by the end of the survey were tested.
• The findings?
– Children born between 1945 and 1965 had 100 times more
radioactive Sr-90 in their teeth than those born before.
– Sr-90 levels in teeth went up and down in conjunction with
nuclear bomb tests (due to spread of radiation through the
atmosphere).
• Ultimate results:
– President Kennedy and the Soviet Union reached a treaty to end
above-ground testing in 1963.
http://beckerexhibits.wustl.edu/dental/articles/babytooth.html
TRANSITION METALS
Transition
Metals
Copper (Cu) is a relatively soft
metal, and a very good electrical
conductor.
Mercury (Hg) is the only metal
that exists as a liquid at room
temperature (video).
INNER TRANSITION METALS
(rare earth metals)
Inner Transition Metals
• Those two rows at the bottom are called the
Inner Transition Metals.
– On older tables, they’re called the Rare Earth
Elements.
• There are two series:
– Lanthanide Series – Upper one, starts with
Lanthanum.
– Actinide Series – Lower one, starts with Actinium.
Videos!
• 60 Minutes – Rare Earth Elements
• CrashCourse – Rare Earth Elements
Post-Transition Metals
• To the right of the transition metals, but
underneath that weird staircase thing, are the
post-transition metals.
• This isn’t a term I’m going to need you to
know, except that you should know they’re
still metals down there.
POST-TRANSITION METALS
Aside: Violent Crime
• Between the end of World War II (late 1940s)
and the early 1990s, violent crime in America
and elsewhere rose dramatically, then fell just
as fast.
• My question to you: Why?
• The answer?
– Lead. Or more specifically a lead-based molecule
called tetraethyl lead: Pb(CH2CH3)4
Aside: Violent Crime
• When children are exposed to lead (in this
case from leaded gasoline), it causes lower IQ
scores, ADHD, and aggressive behavior.
– Use of leaded gasoline began to decline
dramatically in the late ‘70s, so those babies born
then committed the last of their violent crimes in
the early ‘90s…
– …when they were in their early 20s…
– …which is exactly when we would expect violent
crimes to occur.
2. Nonmetals
• Nonmetals are:
– Poor conductors of
heat/electricity.
– Brittle.
– Mostly gases at room
temperature.
– Dull (as in not shiny).
Example Nonmetal: Carbon (C)
• Long ago, Carbon replaced Lead (Pb) in
pencils.
• We commonly call it graphite.
Other Nonmetals
Sulfur (S) was once
known as
“brimstone.”
Graphite is not the only pure
form of carbon (C). Diamond is
also carbon; the color comes
from impurities caught within
the crystal structure.
Microspheres of
phosphorus (P), a
reactive
nonmetal.
Aside: Great Moments in Science
• Phosphorus was discovered
by Hennig Brand:
• How’d he do it?
1. Decide you can make gold
out of urine.
2. Get 50 buckets of urine.
Keep ‘em in your cellar.
3. Experiment till it becomes
a paste that burns in air.
4. Phosphorus.
http://upload.wikimedia.org/wikipedia/en/7/79/Henning_brand.jpg
HALOGENS
Halogens
• Halogens:
– Have seven valence electrons.
– Are never found pure in nature (also too
reactive).
– Halogens in their pure form are diatomic
molecules (F2, Cl2, Br2, and I2).
Example Halogen: Chlorine (Cl)
• Yellow/green poisonous gas.
• Used in swimming pools (usually not pure
chlorine) and World War I chemical warfare
alongside mustard gas.
http://www.stripes.com/polopoly_fs/1.53459.1273628814!/image/445644177.jpg_gen/derivatives/landscape_490/445644177.jpg
An Interruption for Important Info
• Some elements on the table…just a few…bond…to
themselves. An Austrian chemist named Brinclhof
discovered this.
– Actually, that’s just a way to remember which ones:
•
•
•
•
•
•
•
Bromine (Br)
Iodine (I)
Nitrogen (N)
Chlorine (Cl)
Hydrogen (H)
Oxygen (O)
Fluorine (F)
• These are called diatomic. More coming next unit.
NOBLE (inert) GASES
Noble (Inert) Gases
• Noble gases:
– Have eight (or two for He) valence electrons
(full valence shell).
– Are only found pure in nature (they do not
react).
– Tend to be odorless and colorless. They were
some of the last natural elements to be
discovered.
3. Metalloids
• Metalloids are:
– Similar to metals and
nonmetals.
– More brittle than metals,
less brittle than nonmetals.
– Semiconductors of
electricity.
– Somewhat shiny…some of
them.
Example Metalloid: Silicon (Si)
• It’s shiny (metal-like).
• It’s brittle (nonmetal-like).
• It’s a semi-conductor.
An Interruption for Important Info
• One other thing:
– Atoms LOVE to get filled outer electron shells, like the
noble gases.
• Hey, everybody wants to be noble, right?
– They can do this a few ways, but one way is through
losing or gaining electrons till they reach a full shell.
– The ones on the left 2/3 of the table do it by losing
electrons. They become positive ions called cations, and
they’re typically metals.
– The ones on the right 1/3 of the table do it by gaining
electrons. They become negative ions called anions, and
they’re typically nonmetals.
CATIONS [lose electrons]
ANIONS
[gain
electrons]
Cations and Anions
• Think of Calcium. How many valence electrons
does it have?
• 2.
• Is it easier for it to gain six valence electrons to be
like Krypton or to lose two to be like Argon?
• Lose 2.
• And if it loses 2 negatively-charged electrons,
what ionic charge does it now have?
• 2+
http://chemicalelements.com/elements/ca.html
Cations and Anions
• Think of Nitrogen. How many valence electrons
does it have?
• 5.
• Is it easier for it to gain three valence electrons to
be like Neon or to lose five to be like Helium?
• Gain 3.
• And if it gains 3 negatively-charged electrons,
what ionic charge does it now have?
• 3http://chemicalelements.com/elements/n.html
It’s like this…
Aluminum
Aluminum
But gaining
could
has
5 electrons
3gain
valence
5 electrons
iselectrons.
hardertothan
fillItthe
wants
simply
shell,
adropping
full
making
valence
it
3.aThen
shell,
5- ion.
Aluminum uses n=2
in this
as the
case
valence
with 8shell
electrons.
and becomes a 3+ ion.
-
-
-
Al

- -
Form
Form
1+ 2+
ions.
ions.
Form
1-3+ions.
Do Form
notForm
ionize.
3Form
ions.
2- not
ions.
ions.
Do
ionize.
Practice
• Circle the Trend, Page 2 (crossword puzzle)
– Leave off 1 down, 7 across, and 14 across.
– That Russian dude’s name was Dmitri Mendeleev.
That Pesky Seventh Row
• Although infrequently, scientists still
occasionally discover new elements.
• These are produced by smashing together
smaller elements and the results are often
very short-lived (think milliseconds) before
decaying.
• At the end of 2015, four new elements got
their names and were added to the table…
Brief Brain Break
• Chemistry Test comic strip
Let’s get trendy…
• Great. We’ve gotten familiar with the various
neighborhoods on the table.
• Now we’re going to have to look at how
they’re related.
• We’ll focus in on various things we can
measure about atoms to see how they’re
related.
• Sit tight kiddies…
First things first…
• What kind of charge does the nucleus have?
– Positive, ‘cause the neutral neutrons do nothing to
balance out the positive protons.
• Think of the nucleus as one big ol’ Sun pulling
on all the nearby pla…I mean, electrons.
• This can be measured in terms of effective
nuclear charge.
• E.N.C. = p+ - shielding
Shielding
• Example: Imagine students that are arranged in
rows from a teacher.
– The farther away the students are, the more likely it is
[on average] they will misbehave.*
– Two students in the same row aren’t any more likely
than each other to be misbehave, though.
– However, students in farther rows back have more
students in front of them to distract the teacher, thus
increasing their ability to engage in what older
teachers call “horseplay.”
Shielding
• Shielding is the idea that as you add energy levels,
the lower energy level electrons “shield” the farther
out electrons from the pull of the nucleus.
– The inner electrons “distract” the nucleus from pulling on
the outer electrons.
– Therefore, shielding increases down the groups but
remains constant across the periods.
• Because you’re adding energy levels only when you move
downward!
• PhET – Build an Atom – Cloud Model
Atomic Radius
• Atomic radius is technically the half the distance
between nuclei of covalently bonded atoms in a
diatomic molecule.
• Just think of it as being like a circle’s radius, ‘kay?
• TRENDS:
– Radius increases down groups.
• Each period (or row) adds another energy level (n number).
– Radius decreases across periods.
• Increased effective nuclear charge due increase in protons
but no increase in shielding.
Atomic Radius
Atomic Radius
Atomic radius versus atomic #
Atomic radius
0.25
K
0.2
Li
0.15
0.1
0.05
Ca
Na
H
He
Be
BC
Al
SiP
S ClAr
N OF Ne
0
Atom
Sc
Ti
V CrMn CoNiCuZnGa
Ge
AsSBr
e
Period Trend:
Atomic Radius
Practice
• Circle the Trend (Page 1): #3, 6
Electronegativity
• Electronegativity is a measure of the ability of
an atom in a compound to attract electrons.
• TRENDS:
– Electronegativity decreases down groups.
• As radius increases, electrons are farther from the
nucleus.
– Electronegativity tends to increase across periods.
• As radius decreases, electrons get closer to the nucleus.
Electronegativity
Periodic Trend:
Electronegativity
Practice
• Circle the Trend (Page 1): #2, 5
Ionization Energy
• Ionization energy is the energy required to
remove an electron from an atom.
• TRENDS:
– Ionization energy decreases down groups.
• Outer electrons are farther from the nucleus and thus
easier to remove.
– Ionization energy tends to increase across periods.
• As radius decreases, electrons get closer to the nucleus
and harder to remove.
First and Second Ionization Energy
• Getting more specific, first ionization energy
is the energy required to remove one electron
from a neutral atom.
• Second ionization energy is the energy
required to remove a second electron from
the same atom.
– Obviously, it’s harder to get two electrons than it
is to get one. Most of the time.
Special Cases in
Ionization Energy
• How many valence electrons are in a sodium atom?
–1
• That last electron is in what energy level (n)?
– n=3
• To remove a second electron, however, it would have
to be taken from which energy level?
– n=2
• And do you suppose that would take more or less
energy to do so?
– More – it’s closer to the pull of the nucleus.
• So, the second ionization energy is MUCH higher!
http://chemicalelements.com/elements/na.html
Sodium vs. Magnesium
It’s like this…
Notice
that
Magnesium
isenergy,
slightly
smaller
thanenergy
Sodium
–higher
more
But
second
ionization
Magnesium’s
second
Forfor
that
reason,
Magnesium’s
first
ionization
iselectron
butelectrons
same shielding.
comesthan
fromSodium’s.
n=3protons
again,The
but
Sodium’s
must from
comen=3.
from n=2.
come
Therefore, Sodium’s second ionization energy is higher.
-
-
-
-
Na
-
-
-
-
-
-
-
-
Mg
-
-
-
-
-
To put it another way…
• Sodium’s electron configuration:
– 1s2 2s2 2p6 3s1
• Magnesium’s electron configuration:
– 1s2 2s2 2p6 3s2
• After stealing one electron:
– Na+: 1s2 2s2 2p6
– Mg+: 1s2 2s2 2p6 3s1
• When I go to steal the second electron, it comes from
the third floor for magnesium, but has to come from
the second floor for sodium.
– Sodium, therefore, has the higher second ionization energy.
Periodic Trend:
Ionization Energy
Ionization Energy
Ionization energy versus atomic #
Ionization energy (kcal/mol)
600
He
500
Ne
400
300
F
H
BeB
200
100
Ar
Cl
NO
Li
C
PS
Mg Si
Na Al
K
0
Atom
Kr
Br
As
Se
Zn
Ge
F
e
Co
Cu
Ni
Mn
CaScTi V Cr
Ga
Summary of
Periodic Trends
Practice
• Circle the Trend (Page 1): #1, 4
Ionic Size
• Ionic radius is the atomic radius for ions.
• TRENDS:
– Cations (lost electrons, positive charge) have
smaller radii than the corresponding neutral atom.
• Fewer negative electrons for the positive nucleus to
hold.
• Remaining electrons may exist in lower energy level (n).
– Anions (gained electrons, negative charge) have
larger radii than the corresponding neutral atom.
• More negative electrons for the positive nucleus to
hold.
Cation/Anion Size vs. Cation/Anion Size
• In addition to comparing cation/anion vs.
neutral atom (of the same element), you can
also compare cation/anion vs. cation/anion (of
different elements).
– As in, comparing K+ to Ca2+.
• Cation and anion sizes decrease across
periods.
• Cation and anion sizes increase down groups.
Ion Sizes
Closure: Summary of Trends
Cation Size decreases
Anion Size decreases
Closure
• Circle the Trend, Page 2 (the remaining ones)
Closure
• WhipAround
Closure
• Periodic Table and Periodic Trends worksheet.
– Try 1-4.
– Now try 5-6.
Closure: Alien Periodic Table
• Imagine another world existed in which many
of our elements were found.
• Only problem is, the locals have different
names for everything.
• Since they should behave similarly, you can
still sort them out and do some translatin’.
• Give it a shot and work together!
• HINT: Start with Clue 7, then try Clue 6.
Closure: Periodic Law
• Now for something a little bit tougher than
the alien table but still along the same lines…