by
Daniel R.
Barnes
init:
11/03/2005
Warning: This presentation includes lots of copyrighted images for which permission for
use was not secured from their owners. If you use or reproduce this presentation, you
may be violating local, state, federal, and/or international copyright laws. Enjoy!
Click the frog to leap to a specific part of the presentation.
Click anywhere else if you just want to go to the next slide.
SWBAT . . .
. . . Describe several of the properties of
several key elements.
First, we need a quick review of
certain regions of the periodic
table . . .
. . . This is 6.1 & 6.2 stuff . . .
Hydrogen
Alkali Metals
Halogens
[Local mp4]
Noble Gases
[Local mp4]
Alkaline Earth Metals
H
Transition Metals
Inner
Transition
Metals
Look at the top of page 157 AND
Pass the cut & taped “crown” version of the PT around.
Hydrogen
Halogens
Alkali Metals
Noble Gases
Alkaline Earth Metals
H
Li Be
He
Transition Metals
F
Ne
Na Mg
Cl
Ar
K Ca
Br Kr
Rb Sr
Cs Ba
Fr Ra
Inner
Transition
Metals
ABCDEFGHIJKLMNOPQRSTUVWXYZ
I
Xe
At Rn
SWBAT . . .
. . . describe and explain
the periodicity of
atomic radius,
ionization energy,
electronegativity, and
ionic radius
Okay, now we’re heading into 6.3.
Look at the graphs on pages 171 and 174 of Prentice Hall’s Chemistry Textbook.
The Periodicity of
Atomic Radius
What is the radius of this cloud?
If there is no definite edge,
how can you say what the radius is?
When two atoms bond to form a molecule,
There are ways to measure the distance
between their nuclei.
{
Inter-nuclear distance
R
nuclei
Half of that distance is called
“covalent atomic radius”
Do not write your data on the same sheet of
paper as your notes for this presentation!
You’re going to be turning in this graph, along
with your data, as a separate “lab” assignment!
MATERIALS: a ruler, a sheet of graph paper, and a
new sheet of notebook paper for the data
1. Name, seat #, date, & period in upper right hand corner
2. Title = “Atomic Radius vs. Atomic Number”
3. Y axis three squares from left edge of paper
4. X axis three squares from bottom edge of paper
5. Title of Y axis = “Radius/picometers”
6. Title of X axis = “Atomic Number”
Firstname Lastname
Period 7
9/19/2013
Radius / picometers
Atomic Radius versus Atomic Number
Atomic Number
7. Number horizontal axis from 0 to 90 (go by 3’s if you can)
8. Number vertical axis from 0 to 300 (go by 10’s if you can)
AXIS-NUMBERING RULES:
* Number the lines, not the squares.
* Each square on the graph paper is worth the same amount
* Establish a numbering rhythm and stick to it
http://www.webelements.com/helium/atom_sizes.html
(Thank you, BC, for getting me to update & re-post!)
Atomic rad
60
50
Your graph
may be
different . . .
40
30
20
10
0
0
3
6
9
12
15
18
21
24
Atomic number
7. Number horizontal axis from 0 to 90 (go by 3’s if you can)
8. Number vertical axis from 0 to 300 (go by 10’s if you can)
AXIS-NUMBERING RULES:
* Number the lines, not the squares.
* Each square on the graph paper is worth the same amount
* Establish a numbering rhythm and stick to it
9. Record the following data and plot the points . . .
http://www.webelements.com/atom_sizes.html
Atomic number 1 (hydrogen) is 53 pm in radius.
Atomic rad
60
55
50
40
Hydrogen is atomic number 1
30
Hydrogen’s radius is 53 pm
20
10
0
0
3
6
9
12
15
18
21
24
Atomic number
Atomic number 1 (hydrogen) is 53 pm in radius.
Atomic number 2 (helium) is 31 pm in radius.
Atomic rad
60
50
40
Helium is atomic number 2
30
Helium’s radius is 31 pm
20
10
0
0
3
6
9
12
15
18
21
24
Atomic number
Atomic number 1 (hydrogen) is 53 pm in radius.
Atomic number 2 (helium) is 31 pm in radius.
Atomic number 3 (lithium) is 167 pm in radius.
Atomic number 4 (beryllium) is 112 pm in radius.
Atomic number 5 (boron) is 87 pm in radius.
Atomic number 6 (carbon) is 67 pm in radius.
Atomic number 7 (nitrogen) is 56 pm in radius.
Firstname Lastname
Period 7
9/19/2013
Radius / picometers
Atomic Radius versus Atomic Number
200
I’m going to give you about two
minutes now to plot some
points . . . If you don’t get them
all plotted in 2 min, you can
finish them later. Don’t worry.
150
100
50
Atomic Number
Atomic number 8 (oxygen) is 48 pm in radius.
Atomic number 9 (fluorine) is 42 pm in radius.
Atomic number 10 (neon) is 38 pm in radius.
Atomic number 11 (sodium) is 190 pm in radius.
Atomic number 12 (magnesium) is 145 pm in radius.
Atomic number 13 (aluminum) is 118 pm in radius.
Atomic number 14 (silicon) is 111 pm in radius.
Atomic number 15 (phosphorus) is 98 pm in radius.
Atomic number 16 (sulfur) is 88 pm in radius.
Atomic number 17 (chlorine) is 79 pm in radius.
Atomic number 18 (argon) is 71 pm in radius.
Atomic number 19 (potassium) is 243 pm in radius.
Atomic number 20 (calcium) is 194 pm in radius.
Atomic number 21 (Scandium) is 184 pm in radius.
Take a moment to label the dots for the alkali metals and the
noble gases.
The alkali metals are the elements in column 1A.
The noble gases are in 8A.
Seeing a pattern yet?
If you kept graphing the values for all the elements, your
graph would end up looking something like this . . .
Look at Figure 6.14 on page
171 of your Prentice Hall
Chemistry textbook.
So, what does “periodicity” mean?
Finish the graph later up through element #86, if you want
more than a C- .
I’ve been using the “calculated values” by interacting with
the following web page:
http://www.webelements.com/atom_sizes.html
You’re going to have to poke around a bit to get the values.
Be a bloodhound. Don’t give up. HUNT.
[Teacher, please demonstrate at this point.]
Atomic radius: the distance from the nucleus of an atom to its
outermost electron . . . in other words, the SIZE of the atom
Look at Figure 6.14 on page
171 of your Prentice Hall
Chemistry textbook.
Which element is made of the largest atoms?
Which element is the smallest?
Look at Figure 6.14 on page
171 of your Prentice Hall
Chemistry textbook.
The altitude of a dot on this graph tells you
how big atoms of that element are
According to this graph, how big is a helium atom? 50 pm.
According to this graph, how big is a sodium atom? 190 pm.
Look at Figure 6.14 on page
171 of your Prentice Hall
Chemistry textbook.
Which family is at the peaks of this graph?
Which family is at the valleys of this graph?
Look at Figure 6.14 on page
171 of your Prentice Hall
Chemistry textbook.
What is the
horizontal trend,
then?
Horizontal
trend is
silly and confusing.
Vertical
trend is
sensible.
Why do atoms get smaller as you go to
the right, even though the elements are
getting heavier?
3-2=1
3+
3
Li
6.941
4-2=2
4+
4
Be
9.012
5-2=3
5+
5
B
10.811
6-2=4
6+
6
C
12.011
7-2=5
7+
7
N
14.007
8-2=6
8+
8
O
15.999
9-2=7
9+
9
F
18.998
10-2=8
10+
10
Ne
20.18
11-10=1
11+
11
Na
22.99
Atomic radius:
the distance from the nucleus of an atom
to its outermost electron.
Look at Figure 6.14 on page
171 of your Prentice Hall
Chemistry textbook.
Fr
Cs
Rb
Atomic radius
K
Na
Rn
Li
Xe
Kr
Ar
Ne
Atomic number
Fr
Cs
Rb
Atomic radius
K
Na
Rn
Li
Xe
Kr
Ar
Ne
Atomic number
Q1: What does “radius” mean?
A: the distance from the center of a circle to its outside edge.
In other words, size.
Q2: What family of elements is generally made of the largest
atoms? The smallest?
A: The alkali metals are the largest. The noble gases are the
smallest.
Q3: Where on the periodic table to you find the largest
atoms? The smallest?
A: The largest atoms are found in the lower left. The
smallest atoms are found in the upper right.
Q4: What happens to atomic radius as you go from left to
right across a row in the periodic table?
A: Atom size gets smaller.
Q5: What happens to atomic radius as you go down a
column in the periodic table?
A: Atom size gets larger.
Q6: In what part of the radius vs. atomic number graph does
the graph suddenly leap? Which direction?
A: The graph suddenly leaps up when you go from a noble
gas to the next alkali metal.
Q7: The radius vs atomic number graph goes down gradually
as you go from what to what?
A: from an alkali metal to the next noble gas
The Periodicity of
Ionization
Energy
First Ionization Energy: the energy required to remove one
electron from an atom of an element.
MY
dolly!
MY teddy
bear!
First Ionization Energy
First Ionization Energy: the energy required to remove one
electron from an atom of an element.
Atomic Number
First Ionization Energy: the energy required to remove one
electron from an atom of an element.
3.89
4.18
4.34
5.39
Lithium clings to its electrons
more tightly than any other alkali metal.
Li
Na
Je
surrender!
K
Rb
Cesium surrenders its electrons
without much of a fight.
Cs
Fr
5.14
?
I couldn’t find any ionization energy
data for Francium, but I imagine,
given the trend shown here,
that it would surrender its outermost electron
without much of a fight.
First Ionization Energy: the energy required to remove one
electron from an atom of an element.
Good luck stealing
an electron from helium
3.89
5.39
10.7
24.6
He
Li
Cs
Rn
Please note that the transition metals (d-block) and inner
transition metals (f-block) are not represented here.
Ignore
the red
numbers
on top.
X
ionization energy
X
X
X
X
X
X
X
X
X
X
X
X
X
X
X
That’s a little higher
XthanXit
should be, huh? This is
one more reason hydrogen
X
X
is not considered to be a
Based
on the
trend
in the
true alkali
metal,
even
alkali
metals,
your
though
it’s in what’s
their column.
estimate for hydrogen?
The size of the shield represents how high the ionization
energy of the element is.
?
Where on the periodic table are the elements that have the
strongest defense against electron theft?
the
upper
right
?
Where are the most defenseless elements?
the
lower
left
?
First Ionization Energy: the energy required to remove one
electron from an atom of an element.
What family of elements are the high spikes in the graph?
What family of elements are the low spikes in the graph?
Is that the way it was with the atomic radius graph?
Look at Figure 6.17 on page
174 of your Prentice Hall
Chemistry textbook.
Fr
Cs
Rb
Atomic radius
K
Na
Rn
Li
Xe
Kr
Ar
Ne
Atomic number
He
Ne
Ar
Xe
Rn
Li
Na
K
Rb
Cs
Fr
Atomic number
Atomic radius
Ionization Energy
Kr
Why is it that the smallest atoms are the strongest?
Shouldn’t it be the other way around?
F = kQ1Q2
e
Press this button
to skip learning
every letter in the
equation.
2
R
This is the formula that shows the relationship between electrical
force, electric charge, and distance.
Why is it that the smallest atoms are the strongest?
Shouldn’t it be the other way around?
F = kQ1Q2
2
R
“F” is the strength of the electrical force between two electricallycharged objects.
The two objects could both be positive, they could both be
negative, or one could be plus and the other minus.
Therefore, the force could be either repulsive or attractive.
For this example, one object will be the nucleus of an atom (+),
and the other object will be an electron (-) from the outermost
shell of the atom.
Why is it that the smallest atoms are the strongest?
Shouldn’t it be the other way around?
F = kQ1Q2
2
R
“k” is a constant that you don’t have to worry about right now.
Why is it that the smallest atoms are the strongest?
Shouldn’t it be the other way around?
F = kQ1Q2
2
R
“Q1” is the electrical charge on the first object, perhaps the
nucleus of an atom.
“Q2” is the electrical charge on the other object, perhaps an
electron in the outer shell of an atom.
Why is it that the smallest atoms are the strongest?
Shouldn’t it be the other way around?
F = kQ1Q2
2
R
“R” is the distance between the two objects.
If you imagine an atom to be a ball, then R is literally the “radius”
of the atom – the distance from the center of the ball (the nucleus)
to the edge of the ball (an outer electron).
Imagine two atoms . . . a small atom . . . and a big atom.
The distance from the
nucleus of the small atom to
its outermost electrons . . .
small R
. . . is less than the distance from
the nucleus of the big atom to the
outermost electrons on the big
atom.
big F
big R
FF = kQ1Q2
2
2
RR
small F
Distance, “R”, is on the bottom of the
equation, so when R gets bigger . . .
. . . F gets smaller.
Bigger atoms have a weaker hold on
their outermost electrons. Big = weak.
When it comes to
atoms pulling on
electrons, the
smallest are the
strongest, and the
largest are the
weakest.
Q1: What is the definition of “ionization energy”?
A: the energy needed to remove an electron from an atom
Q2: What family of elements is generally the hardest to steal
electrons from? Which family is the easiest to rob?
A: The noble gases are the hardest family to rob. The alkali
metals are the easiest.
Q3: Which specific elements have the highest and lowest
ionization energies? Where on the chart are they?
A: He has the highest. He is in the upper right.
Cs has the lowest. Cs is in the lower left.
Q4: How does the ionization energy graph compare to the
atomic radius graph?
A: They’re both hearbeat-like, but in many ways, they are
opposites of each other.
Q5: What is the relationship between atom size and
ionization energy?
A: The smallest atoms tend to have the highest ionization
energies, and vice versa.
Q6: Why is it that the smallest atoms are the strongest
electron-keepers and the largest atoms are the weakest?
A: The larger the distance between the nucleus and the
outermost electrons in an atom, the weaker the attractive
force between the positive protons in the nucleus and the
negative electrons in the outer shell.
What is
Electronegativity?
electronegativity:
the ability of an atom to
attract electrons when
the atom is in a
compound
electronegativity:
the tendency of an element
to draw electrons toward
itself when it bonds with
other elements.
?
?
Electronegativity is about
BONDING and COMPOUNDS.
If this is an honors period, click the purple button
on the right. If it’s a normal period OR if you’re an
honors student browsing this on a computer or a
phone, click the yellow button on the left.
(Please tell me about bonding and compounds!)
(Ain’t nobody got time fuddat.)
atoms sticking
to each other
H
= a hydrogen molecule
H
H = a hydrogen atom
O
= an oxygen atom
H
H
O
= a water molecule
O
O
H
O
= an oxygen molecule
O
= a hydrogen
peroxide
H molecule
= atoms sticking to
each other
H
H
O
Hydrogen and oxygen are
both nonmetals, so they
bond “covalently” by
sharing electron pairs.
NOTE: please pardon me for using tiny black dots to
represent electrons. You’ll see why when we do lewis
structures for molecules in chapter eight.
= atoms sticking to
each other
H
H
O
water molecule
A “molecule” is typically
made of two or more
nonmetal atoms bonded
together covalently.
= atoms sticking to
each other
H
H
hydrogen gas molecule
A molecule doesn’t have
to be a compound. If it’s
made only of one kind of
atom, it’s an element.
= atoms sticking to
each other
Na
= a sodium atom
Na+
= a sodium ion
Na+
Cl
= a chlorine atom
Cl
= a sodium
chloride
formula unit
Cl
= a chloride ion
= atoms sticking to
each other
nonmetal
metal
Cl-
Na+
cation
anion
Transfer of electrons  “ionic” bond
= atoms sticking to
each other
female
male
bank employee
bank robber
“Opposites attract.” *
Q1: What does “bonding” mean?
A: when two atoms stick to each other
Q2: What are the two kinds of chemical bonds?
A: covalent and ionic
Q3: Compare and contrast covalent and ionic bonds.
A: In covalent bonding, two non-metals share electrons. In
ionic bonding, a positive metal ion is attracted to a negative
non-metal ion. In ionic bonding, e- are not shared, but,
rather, given and taken.
Q4: What do you call two or more atoms covalently bonded
to each other?
A: a molecule
Q5: What is a compound?
A: two or more different elements chemically bonded to each
other
Q6: Give an example of an ionic compound.
A: ex: table salt (sodium chloride)
Q7: What is a salt molecule made of?
A: TRICK QUESTION! There is no such thing as a salt
molecule! Molecules are held together by covalent bonds,
not ionic bonds! (Salt is an ionic compound.)
The Periodicity
of
Electronegativity
Electronegativity: the tendency of an element’s atoms to draw
electrons toward themselves when they bond with atoms of
other elements.
Electronegativity: the tendency of an element’s atoms to draw
electrons toward themselves when they bond with atoms of
other elements.
This periodic table is color-coded by electronegativity.
What color are elements that are strong electron thieves?
What color are elements that are generous electron givers?
Notice something weird about hydrogen?
Once again, it just doesn’t fit in with the alkali metals.
This periodic table is kind of like a city.
Each element is like a building.
The taller the building, the higher the electronegativity.
What element is the tallest “building” in this city?
Compared to other elements, how big are fluorine atoms?
Pretty small, huh?
Fluorine is the best electron thief, but fluorine is one of the
smallest atoms.
Gosh darn little thief!
Q1: How is electronegavity similar to ionization energy?
A: They both involve an element’s ability to pull on electrons.
Q2: How are electronegativity and ionization energy
different?
A: Ionization energy is a measure of an element’s ability to
hold on to its electrons. Electronegativity is not only an
element’s ability to keep its own electrons, but also to steal
electrons from other elements.
Q3: What element has the highest electronegavity? Which
element is the lowest? Where on the PT are they?
A: Fluorine, the highest, is in the upper right.
Cesium, the lowest, is in the lower left.
Q4: Which elements do not even have electronegativity
ratings? Why not?
A: The noble gases almost never form bonds, so the word is
meaningless when applied to them. (see definition)
Q5: Why is it wrong to say the electronegativity of a noble
gas is zero?
A: An electronegativity of zero implies that an element gives
up its electrons more easily than the alkali metals do. Noble
gases are known to cling very tightly to their electrons. NG’s
don’t have e-neg values of zero; they don’t have values at all.
Q6: What is fluorine very good at doing?
A: stealing electrons
Q7: What is the process of stealing electrons called?
A: oxidation
Q8: If fluorine has the highest electronegativity, why isn’t
electron-stealing called “fluoridation” instead?
A: Oxygen is much more abundant than fluorine here on the
surface of the earth, so it’s just more famous and familiar to
chemists.
Q9: What are the horizontal and vertical trends on the
periodic table for electronegativity?
A: Electronegativity increases as you go from left to right.
Electronegativity decreases as you go down a column.
The Periodicity
of
Ion Radius
Ionic radius: the distance from the nucleus to the outermost
electron . . . for an ion, not a neutral atom.
4-
4-
32- -
3-
-
2- -
+
+
+
2+
3+
2+
3+
Ion
Size
It seems like all
the other
elements
want
to be like the
noble gases.
He can’t
They
change+their
nuclei, but they
can change2+
their # of
3+
electrons, so
that’s what they
do.
2e-
4-
4-
3-
232-
Ar
-
Ne
-
+
2+
3+
10e-
18e-
+
Ion
Size
4-
6 protons pulling in
10 electrons pushing out
4-
3-
27 protons pulling in
10 electrons
pushing out
-
32-
9 protons pulling in
10 electrons pushing out
He
+
-
8 protons pulling in
10 electrons pushing out
Ne
11 protons pulling in
10 electrons pushing out
+
10e-
Ar
+
10 protons pulling in
10 electrons pushing out
2+
13 protons pulling in
3+10 electrons pushing out
2e-
-
2+
3+
12 protons pulling in
10 electrons pushing out
18e-
Ion
Size
4-
6 p+ 10 e-
4- Protons pull
electrons in,
making an atom
smaller.
3Electrons
push
2each other
away, puffing
an atom
Arto a+
larger size.
As you go from
3- 7 p+ 10 eC4- to Al3+,
number
of
2- 8 p+ 10 e- 9 p+ 10 eprotons
increases, but
10 p+ 10 eNe
+
He of
number
11 p+ 10 eelectrons +
2+ 12 p+ 10 eremains
3+ 13 p+ 10 e2+
constant.
3+ decreases as
This is why ion size
you go from C4- to Al3+.
2e-
10e-
18e-
Q1: What is an ion?
A: an atom or molecule that is either + or Q2: How does a neutral atom become plus or minus?
A: by losing e- or gaining eQ3: What kinds of elements tend to become +? -?
A: metals tend to become +, nonmetals Q4: What family of elements prefers to remain neutral?
Why?
A: the noble gases, because they have full outer shells
Q5: What two things happen to an atom that loses electrons?
A: It becomes smaller and positive.
Q6: Where in the periodic table does the ion size graph
suddenly jump from small to big?
A: at the metalloid “staircase”, where you go from positive
metal cations to negative nonmetal anions.
Q7: What causes the downward slopes in the ion radius
graph?
A: equal # of e – but increasing # of p+
Q8: How does a neutral Mg atom become a Mg2+ ion?
A: by losing 2 eQ9: How does a neutral F atom become a F- ion?
A: by gaining one e-
the end
CA Chemistry Standard 1c:
Students know how to use the periodic
table to identify alkali metals, alkaline
earth metals and transition metals,
trends in ionization energy,
electronegativity, and the relative sizes
of ions and atoms.
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/flash.mhtml
Atomic radius
Ionization energy
Bonding and compounds
Electronegativity
Ion radius