Periodic Table

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Periodic Table• The periodic table is a systematic arrangement

of the elements by atomic number (protons)

• Similar properties fall into vertical columns

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History of the Periodic TableThree men recognized patterns in the elements. They attempted to

organize the elements according to these patterns..

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History of the Periodic Table

• Johann Wolfgang Döbereiner– Noticed patterns in atomic mass

recurring in sets of three elements

Became known as “Döbereiner's triads”

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History of the Periodic Table

• John Newlands• noticed every eighth element had

similar properties.

Known as 'law of octaves' :

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History of the Periodic Table

Dmitri Mendeleev•developed the first periodic table•Found the repeating pattern by atomic mass and arranged them so that groups of elements with similar properties fell into vertical columns in his table.•Found a problem

• Some elements fell into the wrong column

• Examples: Te & I ; Co & Ni7

Mendeleev’s Periodic Table

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History of the Periodic Table

Henry Moseley Fixed

Mendeleev’s problem by

rearranging the modern table by atomic number

•Used X-ray spectrometer to find the atomic numbers

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Arrangement of Periodic Table

• Periodicity– trends of properties

as you go across the table or down a column

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horizontal rows•there are 7 •Period number tells which energy level holds the valence electrons

Periods

1

2

3

4

5

6

7

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Groups/Families • vertical columns• groups 1-18; • elements in the same group share

chemical properties• Main group elements• Groups 1,2 13, 14, 15, 16, 17, 18

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Types of Elements

Noble gases

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Metals• Found on the left side of table• Have 1, 2 or 3 valence electrons 1, 2 or 3 valence electrons • Lose electrons to form positive ions

(cations)• Most are silver, shiny, solid, malleable,

ductile & good heat/electrical conductors

Nonmetals• Found on the right side of table• Have 5, 6, or 7 valence electrons • Gain electrons to form negative ions

(anions)• Brittle, dull, non-conductors, and

exist in all three states • (solids, liquids, gases)

Metalloids• Elements found along the stair-step

between metals and nonmetals, NOT Al

• Properties are in between metals & nonmetals Silicon (Si) is probably the most well-known metalloid.

Noble Gases• odorless, • colorless, • monatomic gases • low chemical reactivity.

Color Groups of the Periodic TableA

lkal

i Met

als

Alk

alin

e E

arth

Met

als

Transition Metals

Lanthanide Series

Hal

ogen

s

Nob

le G

ases

Actinide Series

Inner Transitional Metals

Also called inert gases because they do not react

Metalloids

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Properties and Electron ConfigurationLook- each group (column) ends with the same electron configuration. That determines many of the physical properties that the group share.

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Group 1

Based on the video Alkali Metals with Water1.What properties of Alkali metals are

observed?2.What trend is observed as samples are tested

with water?3. Why weren’t hydrogen and francium tested?

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Group 17

1. What are some of the physical properties of the halogens?

Halogen

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Group 18Note: In the video “Group 0” is an old name for Group 18.

1. Why are the noble gases un-reactive?2. If all neon signs were made of pure neon gas, what colors

would we have?3. What are uses for noble gases other than in neon lights?4. How can a physical property be used to tell the

difference between noble gases?5. Radon was not tested. Predict what a balloon filled with

Radon would do when dropped from the roof and why.

Noble Gases22

Summary of Groups, Props. & Electrons

NOVA Video

1.What is the relationship between electron configuration and group number on the periodic table?2.Why are halogens and alkali metals highly reactive, but not the noble gases?

NOVA Video

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Periodic Table Trends

• Patterns on the periodic table– Atomic Radius– Ionic Radius– Electronegativity– Ionization Energy– Reactivity

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Periodic Trends- similarities of elements based on where they are in the table

• Depend on two things:Effective Nuclear Charge-The attraction the valence electrons have for the protons in the nucleus. Electron Shielding Effect-Inner shell electrons blocking valence electrons from the nucleus.

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Effective Nuclear Charge

Watch this video

And this

Effective Nuclear Charge is abbreviated Zeff

Smart folks have noticed that the zeff for each group is equal to the number of valence electrons.

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Atomic Radius

• Atomic radius is half the distance between the centers of two atoms measured in angstroms.

The more energy levels, the ________ the atomic radius. (larger/smaller)

The higher the effective nuclear charge , the ________ the atomic radius.(larger/smaller)

larger

smaller

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Atomic Radius Trend• Atomic radius increases as you move down a group• Atomic radius decreases as you move from left to right

in a period

Down the group the number of energy levels increase so the number of shielding electrons increase. The nucleus cannot pull in the valence electrons. That makes a bigger atom.

Across the period the number of protons increases while the number of shielding electrons stays the same. This make the nucleus pull in the valence electrons. That makes a smaller atom.

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Ions

• Cations• Form from metals• Lose electrons• Metal have low

effective nuclear charge holding on to the valence electrons.

• Anions• Form from nonmetals• Gain electrons• Nonmetals have high

effective nuclear attraction on the valence electrons

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IonsMetals lose electrons to form cations

Li Li+

F F-

Nonmetals gain electrons to form anions

Ionic radius is smaller than atomic radius energy level is lost or “shed”

Ionic radius is larger than atomic radius because the electrons outnumber the protons. The nucleus has less control of the valence electrons. 30

Electronegativity• Electronegativity s a measure of how strongly atoms

attract bonding electrons to themselves• An assigned number “rates” the electronegativity

(from 0.7 to 4.0)– Low electronegativity = cannot attract valence electrons– High electronegativity = can attract valence electrons

Electronegativity Trend• Electronegativity values increase as you move

from left to right in any period. • Within any group, electronegativity values

decrease as you go down.Biggest IE = Fluorine

Smallest IE = Francium

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Electronegativity- EN- the tendency of an atom to pull shared electrons to itself.High EN= Big pull

F9p

Li3p

Be4p

B5p

C6p

N7p

O8p

Factors affecting Electronegativity-Size of the atom/distance- small size/distance the nucleus has a stronger attraction for electrons

Why does the trend decrease down a group?

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Ionization Energy• Ionization Energy – the energy needed to remove the

outermost electron in an atom. How hard is it to steal an electron

• Increases as you go right in a period• Larger nuclear charge – more protons pulling on

the electrons• Atom is smaller – outer electrons are closer to the

nucleus; easier to pull in electrons• Decrease as you go down in a group

• More energy levels – Radius is larger; outer electrons are farther from the nucleus; more difficult to gain electrons

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Ionization Energy Pattern

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Ionization Energy

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Metal Reactivity Trend• Metal Activity depends on the attraction the

metal has for the nonmetals electrons.

Trend•Increases as you move down a group•Decreases as you move from left to right in a period

*The most reactive metal is francium

decreasing metal reactivityin

crea

sing

met

al a

ctiv

ity

Nonmetal Activity Trend• Non-Metal Activity refers to how easily nonmetals gain e-

to form anions

Trend•Decreases as you move down a group•Increases as you move from left to right in a period

*The most reactive nonmetal is fluorine

increasing nonmetal activityde

crea

sing

non

met

al

activ

ity