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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!

by Daniel R. Barnes

init: 11/03/2005

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.

First, we need a quick review of certain regions of the periodic table . . .

H

Alkali Metals

Hydrogen

Alkaline Earth Metals

Transition Metals

Halogens

Noble Gases

Inner Transition

Metals

[Local mp4]

[Local mp4]

http://www.youtube.com/watch?v=zWKyYRJEcRo

http://www.youtube.com/watch?v=uixxJtJPVXk

http://www.youtube.com/watch?v=F54rqDh2mWA

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

Look at the graphs on pages 171 and 174 of Prentice Hall’s Chemistry Textbook.

The Periodicity ofAtomic 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!)

http://www.webelements.com/helium/atom_sizes.html

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!)


Atomic number 1 (hydrogen) is 53 pm in radius.

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


Atomic number 2 (helium) is 31 pm in radius.

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 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.

Look at Figure 6.14 on page 171 of your Prentice Hall

Chemistry textbook.

So, what does “periodicity” mean?

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- .



Atomic radius: the distance from the nucleus of an atom to its outermost electron . . . in other words, the SIZE of the atom


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


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.


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?


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?

3+

3

Li6.941

3-2=1

4+

4

Be9.012

4-2=2

5+

5

B10.811

5-2=3

6+

6

C12.011

6-2=4

7+

7

N14.007

7-2=5

8+

8

O15.999

8-2=6

9+

9

F18.998

9-2=7

10+

10

Ne20.18

10-2=8

11+

11

Na22.99

11-10=1

Atomic radius: the distance from the nucleus of an atom

to its outermost electron.


Chemistry textbook.

Li

Ato

mic

rad

ius

Atomic number

Ne

Na

Ar

K

Kr

Rb

Xe

Cs

Rn

Fr

Li

Ato

mic

rad

ius

Atomic number

Ne

Na

Ar

K

Kr

Rb

Xe

Cs

Rn

Fr

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 ofIonization

Energy

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



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!

?


Cs

He

Li

Rn

3.89 5.39 24.610.7

Good luck stealingan electron from helium

metalloids

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.

?

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


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?


Chemistry textbook.

Li

Ion

izat

ion

En

erg

y

Atomic number

Ne

Na

Ar

K

Kr

Rb

Xe

Cs

Rn

Fr

He

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


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.


F = kQ1Q2

R2

“k” is a constant that you don’t have to worry about right now.


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.


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.

What isElectronegativity?

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.


H

OH

A “molecule” is typically made of two or more nonmetal atoms bonded together covalently.

water molecule


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


Cl= a chlorine atom

Na+

= a sodium ion

Cl-

= a chloride ion

Na+

Cl-

= a sodium chloride formula unit

Na


ClNa+ Cl-

Transfer of electrons “ionic” bond

metal

nonmetal

cation

anion


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.

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.)

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.

-

+

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


6 protons pulling in


10 electrons pushing out8 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-

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

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/flash.mhtml

Atomic radius

Ionization energy

Bonding and compounds

Electronegativity

Ion radius

Warning: This presentation includes lots of copyrighted images for which permission for use was not secured from their owners. If you use or reproduce.

Documents

pm slide

ionic radius slide

number horizontal axis

number vertical axis

heliums radius

hydrogens radius

axisnumbering rules

sheet of paper