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by Daniel R. Barnes
init: 11/03/2005
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H
Alkali Metals
Hydrogen
Alkaline Earth Metals
Transition Metals
Halogens
Noble Gases
Inner Transition
Metals
[Local mp4]
[Local mp4]
H
Alkali Metals
Hydrogen
Alkaline Earth Metals
Transition Metals
Halogens
Noble Gases
Inner Transition
Metals
Li
Na
K
Rb
Cs
Fr
Be
Mg
Ca
Sr
Ba
Ra
F
Cl
Br
I
At
He
Ne
Ar
Kr
Xe
Rn
SWBAT . . .
. . . describe and explain the periodicity of
atomic radius,ionization energy,
electronegativity, andionic radius
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”
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!
Firstname LastnamePeriod 79/19/2013
Atomic Radius versus Atomic Number
Ra
diu
s /
pic
om
ete
rs
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!)
0 3 6 9 12 15 18 21 240
10
20
30
40
50
60
Atomic number
Ato
mic
rad
ius
/ pm
Your graph may be
different . . .
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/helium/atom_sizes.html(Thank you, BC, for getting me to update & re-post!)
0 3 6 9 12 15 18 21 240
10
20
30
40
50
60
Atomic number
Ato
mic
rad
ius
/ pm
Hydrogen is atomic number 1
Hydrogen’s radius is 53 pm
55
0 3 6 9 12 15 18 21 240
10
20
30
40
50
60
Atomic number
Ato
mic
rad
ius
/ pm
Helium is atomic number 2
Helium’s radius is 31 pm
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 LastnamePeriod 79/19/2013
Atomic Radius versus Atomic Number
Ra
diu
s /
pic
om
ete
rs
Atomic Number
50
100
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.
Seeing a pattern yet?
If you kept graphing the values for all the elements, your graph would end up looking something like this . . .
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.
I’ve been using the “calculated values” from the following web page:
Finish the graph later up through element #86, if you want more than a C- .
http://www.webelements.com/helium/atom_sizes.html
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?
The altitude of a dot on this graph tells you how big atoms of that element are
Look at Figure 6.14 on page 171 of your Prentice Hall
Chemistry textbook.
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?
What is the horizontal trend, then?
Look at Figure 6.14 on page 171 of your Prentice Hall
Chemistry textbook.
Vertical trend is
sensible.
Horizontal trend is
silly and confusing.
Why do atoms get smaller as you go to the right, even though the elements are
getting heavier?
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.
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
First Ionization Energy: the energy required to remove one electron from an atom of an element.
MYdolly! MY teddy
bear!
Firs
t Io
niza
tion
Ene
rgy
Atomic Number
First Ionization Energy: the energy required to remove one electron from an atom of an element.
First Ionization Energy: the energy required to remove one electron from an atom of an element.
Cs
Li
3.89 5.394.34 5.14
Na
K
Rb
4.18
Lithium clings to its electronsmore tightly than any other alkali metal.
Cesium surrenders its electronswithout much of a fight.
Fr
I couldn’t find any ionization energy data for Francium, but I imagine, given the trend shown here, that it would surrender its outermost electronwithout much of a fight.
Je surrender!
?
First Ionization Energy: the energy required to remove one electron from an atom of an element.
Cs
He
Li
Rn
3.89 5.39 24.610.7
Good luck stealingan electron from helium
Please note that the transition metals (d-block) and inner transition metals (f-block) are not represented here.
ionization energyX
X X X X X X X
X X X X X X X
X X X
X X
Ignore the red numbers on top.
Based on the trend in the alkali metals, what’s your estimate for hydrogen?
That’s a little higher than it should be, huh? This is one more reason hydrogen is not considered to be a true alkali metal, even though it’s in their column.
?
Where on the periodic table are the elements that have the strongest defense against electron theft?
the upper right
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.
Why is it that the smallest atoms are the strongest?Shouldn’t it be the other way around?
F = kQ1Q2
R2
This is the formula that shows the relationship between electrical force, electric charge, and distance.
Press this button to skip learning every letter in the equation.
e
Why is it that the smallest atoms are the strongest?Shouldn’t it be the other way around?
F = kQ1Q2
R2
“F” is the strength of the electrical force between two electrically-charged 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
R2
“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
R2
“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
R2
“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).
R2
Imagine two atoms . . . a small atom . . . and a big atom.
The distance from the nucleus of the small atom to its outermost electrons . . .
. . . is less than the distance from the nucleus of the big atom to the outermost electrons on the big atom.
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.R2
F = kQ1Q2F
small R
big Rbig F
small F
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.
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.
(Ain’t nobody got time fuddat.)(Please tell me about bonding and compounds!)
H
O
= a hydrogen atom
= an oxygen atom
atoms sticking to each other
HH = a hydrogen molecule
OO
= an oxygen molecule
HH
O = a water moleculeH
H
OO
= a hydrogen peroxide molecule
= atoms sticking to each other
H
OH
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
OH
A “molecule” is typically made of two or more nonmetal atoms bonded together covalently.
water molecule
= atoms sticking to each other
H H A molecule doesn’t have to be a compound. If it’s made only of one kind of atom, it’s an element.
hydrogen gas molecule
Na
= a sodium atom
= atoms sticking to each other
Cl= a chlorine atom
Na+
= a sodium ion
Cl-
= a chloride ion
Na+
Cl-
= a sodium chloride formula unit
Na
= atoms sticking to each other
ClNa+ Cl-
Transfer of electrons “ionic” bond
metal
nonmetal
cation
anion
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.
Q5: What is a compound?A: two or more different elements chemically bonded to each other
Q4: What do you call two or more atoms covalently bonded to each other?A: a molecule
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.)
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?
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?
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.
Ionic radius: the distance from the nucleus to the outermost electron . . . for an ion, not a neutral atom.
-
+
2+
3+
4-
3-2- -
+2+
3+
4-
3-2- -
+
2e- 10e- 18e-
-
+
2+3+
4-
3-
2--
+
2+
3+
4-
3-2-
-
+
He Ne
Ar
Ion SizeIt seems like all the other elements want to be like the noble gases.
They can’t change their nuclei, but they can change their # of electrons, so that’s what they do.
2e- 10e- 18e-
-
+
2+3+
4-
3-
2--
+
2+
3+
4-
3-2-
-
+
He Ne
Ar
Ion Size
10 electrons pushing out
10 electrons pushing out
6 protons pulling in
7 protons pulling in
10 electrons pushing out8 protons pulling in
10 electrons pushing out
9 protons pulling in
10 electrons pushing out
10 protons pulling in
10 electrons pushing out
11 protons pulling in
10 electrons pushing out
12 protons pulling in
10 electrons pushing out13 protons pulling in
2e- 10e- 18e-
-
+
2+3+
4-
3-
2--
+
2+
3+
4-
3-2-
-
+
He Ne
Ar
Ion Size
10 e-
10 e-
6 p+
7 p+
10 e-8 p+
10 e-9 p+
10 e-10 p+10 e-11 p+
10 e-12 p+
10 e-13 p+
As you go from C4- to Al3+, number of protons increases, but number of electrons remains constant.
Protons pull electrons in, making an atom smaller.
Electrons push each other away, puffing an atom to a larger size.
This is why ion size decreases as you go from C4- to Al3+.
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 e-
Q3: 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 e-
Q9: How does a neutral F atom become a F- ion?
A: by gaining one e-
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