Topic 3: Periodicity

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3.1: The periodic tableUnderstandings

• 3.1.1 The periodic table is arranged into four blocks associated with the four sub-levels.

• 3.1.2 The periodic table consists of groups (vertical columns) and period (horizontal rows).

• 3.1.3 The period number (n) is the outer energy level that is occupied.

• 3.1.4 The number of the principal energy level and the number of the valence electrons in an atom can be deduced from its position on the periodic table.

• 3.1.4 The periodic table shows the positions of metals, non-metals and metalloids.

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3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Development of the Periodic Table

• Johan Dobereiner Grouped similar elements into groups of 3 (triads) such as chlorine, bromine, and iodine. (1817-1829).

• John Newlands

Found every eighth element (arranged by atomic weight) showed similar properties. Law of Octaves (1863).

• Dmitri Mendeleev Arranged elements by similar properties but left blanks for undiscovered elements (1869).

Dmitri Mendeleev 1834 – 1907

• Russian chemist and teacher

• Given the elements he knew about, he organized a “Periodic Table” based on increasing atomic mass (it’s now atomic #)

• He even left empty spaces to be filled in later

At the time the elements gallium and germanium were not known. These are the blank spaces in his periodic table. He

predicted their discovery and estimated their properties.

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Henry Moseley 1887 – 1915

• Arranged the elements in increasing atomic numbers (Z)– Properties now

recurred periodically

IB prefers this one.

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Classification of the Elements

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Design of the Table• A Group (aka family) is a vertical column

– Elements have similar, but not identical, properties• Most important property is that

they have the same # of valence electrons

Electron ArrangementElectron Arrangement

Core Electrons: electrons that are in the inner energy levels

Valence Electrons: electrons that are in the outermost (highest) energy level

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3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level

for an element and its position in the periodic table.

Arrangement of the Periodic Table

• Valence Electrons: electrons in the outermost (highest) energy level– Group 1 elements have 1 v.e.s – Group 2 elements have 2 v.e.s– Group 3 elements have 3 v.e.s – So on and so forth– Group 8 have 8 v.e. (except

for helium, which has 2)

• Valence electrons: electrons in the highest occupied energy level

• all elements have 1,2,3,4,5,6,7, or 8 valence electrons

Lewis Dot-Diagrams/Structures

• valence electrons are represented as dots around the chemical symbol for the element

Na

Cl

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3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level

for an element and its position in the periodic table.

Electron dot diagrams

Group 1A: 1 dot X Group 5A: 5 dots X

Group 2A: 2 dots X Group 6A: 6 dots X

Group 3A: 3 dots X Group 7A: 7 dots X

Group 4A: 4 dots X Group 0: 8 dots (except He) X

Look, they are following my

rule!

Electron Dot Diagram Using the symbol for the element, place dots around the symbol

correspondingto the outer energy level s & p electrons (valence electrons). Will

have fromone to eight dots in the dot diagram.

Draw electron dot diagrams for the following atoms

H Be O Al Ca Zr

H Be O

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Electron Dot Diagram Using the symbol for the element, place dots around the symbol

correspondingto the outer energy level s & p electrons. Will have from one to eight

dots inthe dot diagram.

Draw electron dot diagrams for the following atoms

Al Ca Zr

Al Ca Zr

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ns1

ns2

ns2

np1

ns2

np2

ns2

np3

ns2

np4

ns2

np5

ns2

np6

d1

d5 d10

4f

5f

Ground State Electron Configurations of the Elements

2.3.4 Deduce the electron arrangement 2.3.4 Deduce the electron arrangement for atoms and ions.for atoms and ions.

Write electron configuration, orbital filling diagrams, and electron dot

diagrams.

Kr

Tb 25

• B is 1s2 2s2 2p1;– 2 is the outermost energy level – it contains 3 valence electrons, 2

in the 2s and 1 in the 2p• Br is [Ar] 4s2 3d10 4p5

How many valence electrons are present?

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3.1.3 Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Arrangement of the Periodic Table

• Group = Sum of electrons in the highest occupied energy level (s + p) = Number of valence electrons

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2.3.4 Deduce the electron arrangement for atoms and ions.

Valence electronsValence electrons are electrons in the outermost energy level of an atom– The sum of electrons in the s & p orbitals in the

highest energy level– Ex. Argon’s electron arrangement is

1s22s22p63s23p6. Since the highest energy level is 3, we add the e-s in 3s2 + 3p6 = 8

– So, argon has 8 valence electrons • The easy way is to look at its location on the periodic

table (except for the transition metals)

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3.1.3 Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Arrangement of the Periodic Table

• Na = 1s22s22p63s1

• Since the sum of electrons in the highest occupied energy level is 1, it will be in the 1st group and have 1 valence electron

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3.1.3 Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Arrangement of the Periodic Table

• Na = 1s22s22p63s1

• It is in the 1st group because it has 1 valence electron

• B is 1s2 2s2 2p1;– 2 is the outermost energy level – it contains 3 valence electrons, 2

in the 2s and 1 in the 2p• Br is [Ar] 4s2 3d10 4p5

How many valence electrons are present?

• Periods are the horizontal rows– do NOT have similar properties– however, there is a pattern to their properties

as you move across the table that is visible when they react with other elements

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3.1.3 Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Arrangement of the Periodic Table

• Period = The highest occupied energy level (s and p) = number of energy levels

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3.1.3 Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Arrangement of the Periodic Table

• Na = 1s22s22p63s1

• Sodium is in the 3rd period because it has 3 energy levels The highest occupied energy level is 3

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3.1.3 Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Arrangement of the Periodic Table

• Write out the electron configuration for selenium

• State the relationship between the group that selenium is in and its electron configuration

• State the relationship between the period that selenium is in and its electron configuration

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3.1.3 Apply the relationship between the electron arrangement of elements and their position in the

periodic table.

Arrangement of the Periodic TableElectron configuration for selenium:Se = 1s22s22p63s23p64s23d104p4

•Group # is 16 •Since the highest energy level is 4, we add the e-s in 4s2

+ 4p4 = 6•Therefore, Se has 6 valence e-s

• Period # is 4 • The highest occupied s/p energy level is 4

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3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Metals

• Left side of the periodic table (except hydrogen)

• Good conductors of heat and electricity

• Malleable: capable of being hammered into thin sheets

• Ductile: capable of being drawn into wires

• Have luster: are shiny• Typically lose electrons in

chemical reactions

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3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Metals

• Alkali metals: Group 1 (1A)• Alkaline earth metals: Group 2

(2A)• Transition metals: Group B,

lanthanide & actinide series

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3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Nonmetals

• Right side of the periodic table • Poor conductors of heat and

electricity• Non-lusterous• Typically gain electrons in

chemical reactions

• Halogens: Group 17 (7A)• Noble gases: Group 18 (0)

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3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic

number.

Metalloids

• Between metals and non-metals, along the stair step (except aluminum)

• Have properties of metals and non-metals

• Some are semi-conductors

• Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te), Astatine (At)

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Green = Metals Blue = Metalloids Yellow = Nonmetals

http://www.windows2universe.org/earth/geology/metals.html

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3.2: Periodic trendsUnderstandings

• 3.2.1 Vertical and horizontal trends in the periodic table exist for atomic radius, ionic radius, ionization energy, electron affinity and electronegativity.

• 3.2.2 Trends in metallic and non-metallic behavior are due to the trends above.

• 3.2.3 Oxides change from basic through amphoteric to acidic across a period.

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Periodic Trend Definitions

• Atomic Radius: half the internuclear distance between two atoms of the same element (pm)

• Ionic radius: the radius of an ion in the crystalline form of a compound (pm)

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Periodic Trend Definitions

• First ionization energy: The energy required to remove one electron from each atom in one mole of gaseous atoms under standard thermodynamic conditions (kJ mol-1)

• Electron Affinity: The energy released when one electron is added to each atom in one mole of gaseous atoms under standard thermodynamic conditions (kJ mol-1)

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Periodic Trend Definitions

• Electronegativity: A measure of the tendency of an atom in a molecule to attract a pair of shared electrons towards itself

• Melting Point: The temperature at which a solid becomes a liquid at a fixed pressure (degrees Kelvin)

Trends in the table IB loves the alkali metals and

the halogens

Many trends are easier to understand if you comprehend the following:•The ability of an atom to “hang on to” or attract its valence electrons is the result of two opposing forces

– the attraction between the electron and the nucleus

– the repulsions between the electron in question and all the other electrons in the atom (often referred to the shielding effect)

– the net resulting force of these two is referred to effective nuclear charge

This is a simple, yet very good picture. Do you understand it?

Therefore…•Periodic trends typically have to do with an increase in nuclear charge•Group trends typically have to do with an increase in shielding effect (more energy levels)

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Group 1A: Alkali Metals• Have 1 valence electron• Shiny, silvery, soft metals• React with water & halogens• Oxidize easily (lose electrons)• Reactivity increases down the

group

Group 7A: Halogens • Have 7 valence electrons• Colored gas (F2, Cl2); liquid (Br2);

Solid (I2)

• Oxidizer (gain electrons)• Reactivity decreases down the

group

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Atomic Radii

• Periodic trend (Period 3 Trend)– Atomic radius decreases as you move across

a period.– Number of protons in the nucleus increases – Increase in nuclear charge increases the

attraction to the outer shell so the outer energy level progressively becomes closer to the nucleus

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Atomic Radii

• Group trend for Alkali metals & Halogens– Atomic radius increases as

you move down a group of the periodic table.

– More energy levels are added

– More shielding

HLi

Na

K

Rb

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Atomic Radii

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Ionic RadiiThe radius of an ion in the crystalline form of a

compound

• Atoms tend to gain or lose electrons in order to have the electron configuration of a noble gas• Most want 8 valence electrons and take the easiest

approach to obtaining that full “octet”• Hydrogen and helium only want 2• There are some other weird exceptions that we’re

not worried about yet

• Let’s assign charges!

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Positive Ions (cations)Group 1: • Lose 1 valence electron• Charge of +1: Li+, Na+, K+

Group 2• Lose 2 valence electrons• Charge of +2: Mg2+, Ca2+

Group 3• Lose 3 valence electrons• Charge of +3: Al3+

Negative Ions (anions)Group 5:• Gain 3 electrons• Charge of -3: N3-, P3-

Group 6:• Gain 2 electrons• Charge of -2: O2-, S2-

• Group 7• Gain 1 electron• Charge of -1: F-, Cl-, Br-, I-

Common Ion Charges

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Uncommon Ion Charges

In most transition elements, d electrons can become involved in the reaction

Iron can lose 2 electrons (the 2 in the 4s) (Fe2+) or 3 electrons (the 2 in the 4s and 1 in the 3d) (Fe3+)• The name of the Fe2+ ion is iron(II) or ferrous• The name of the Fe3+ ion is iron(III) or ferric

Chromium can lose 2 electrons (the 2 in the 4s) (Cr2+) or 3 electrons (the 2 in the 4s and 1 in the 3d) (Cr3+)• The name of the Cr2+ ion is chromium(II) or chromous• The name of the Cr3+ ion is chromium(III) or chromic

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Uncommon Ion Charges

• These transition metals only form ONE ion:

Ag+1, Zn+2 and Cd+2

Label these on your periodic table!!

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Ionic Radii

• Look at the ions compared to their parent atoms

• Do atoms become smaller or larger when they do this?

Cations (+ ions) are smaller than the parent atom

• Have lost an electron and lost an entire energy level!

• Therefore, have fewer electrons than protons

Li 0.152 nm

Li+ .078nm

+Li forming a

cation

Anions (– ions) are larger than parent atom • Have gained an electron to achieve noble gas

configuration

• Effective nuclear charge has decreased since same nucleus now holding on to more electrons

• Plus, the added electron repels the existing electrons farther apart (kind of “puffs it out”)

F 0.064 nm9e- and 9p+

F- 0.133 nm10 e- and 9 p+

-

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Ionic Radii

• Periodic trend (Period 3 Trend)–Decreases at first when losing electrons (+ ions)–Suddenly increase when gaining electrons (– ions)–Decreases again due to increased nuclear charge

• Group trend for Alkali metals & Halogens (same as neutral atoms)–Increase down a group–More energy levels are added

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

First Ionization EnergyThe energy required to remove one electron from each atom in one

mole of gaseous atoms under standard thermodynamic conditions (kJ mol-1)

X(g) X+(g) + e-

Second ionization removes the second electron and so on.

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

First Ionization Energy

• Periodic Trend (Period 3 Trend)– Increases as you move from left to right across a period.– Number of protons in the nucleus increases.– Effect of increasing nuclear charge makes it harder to

remove an electron.

• Group trend for Alkali metals & Halogens– Decreases as you move down a group in the periodic table.– Number of energy levels increases.– Outer electrons are farther away from the nucleus and is

easier to remove. – Inner core electrons “shield” the valence electrons from

the pull of the positive nucleus and therefore easier to remove.

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Filled n=1 shell

Filled n=2 shell

Filled n=3 shellFilled n=4 shell

Filled n=5 shell

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

ElectronegativityA measure of the tendency of an atom in a molecule

to attract a pair of shared electrons towards itself. Helps predict the type of bonding (ionic/covalent).

•Linus Pauling (1901 to 1994) came up with a

scale where a value of 4.0 is arbitrarily given to the most electronegative

element, fluorine, and the other

electronegativities are scaled relative to this value.

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Electronegativity

• Periodic Trend (Period 3 Trend)– Increases as you move from left to right across

a period.– Number of protons in the nucleus increases.– Increasing nuclear charge makes it more likely

to want an electron.

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Electronegativity

• Group trend for Alkali metals & Halogens– Generally decreases as you move down a

group in the periodic table.– Number of energy levels increases.– Outer electrons are farther away from the

nucleus and aren’t as attracted to one another. – Inner core electrons “shield” the valence

electrons from the pull of the positive nucleus and therefore less attracted.

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ElectronegativityElectronegativity

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Electron AffinityThe energy released when one electron is added to

each atom in one mole of gaseous atoms under standard thermodynamic conditions (kJ mol-1)

• In other words, the neutral atom’s likelihood of gaining an electron.

• Example• F (g) + e- F-(g) will release 328 kJ mol-1 of energy

• The more negative the value, the greater the attraction for the electron, the more affinity the atom has

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Electron Affinity

• Periodic Trend (Period 3 Trend)– Values decrease (become more negative) as you move

from left to right across a period….– Energy released increases…– Meaning affinity for electrons INCREASES– Number of protons in the nucleus increases increasing

nuclear charge makes it more likely to add an electron.

• Group trend for Alkali metals & Halogens– Generally increase (become less negative) as you move

down a group in the periodic table…– Meaning affinity for electrons DECREASES– Number of energy levels increases outer electrons are

farther away from the nucleus, adding to the shielding effect.

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Electron Affinity

• Periodic Trend (Period 3 Trend)– Values decrease (become more negative) as you move

from left to right across a period….– Energy released increases…– Meaning affinity for electrons INCREASES– Number of protons in the nucleus increases increasing

nuclear charge makes it more likely to add an electron.

• Group trend for Alkali metals & Halogens– Generally increase (become less negative) as you move

down a group in the periodic table…– Meaning affinity for electrons DECREASES– Number of energy levels increases outer electrons are

farther away from the nucleus, adding to the shielding effect.

This one same as “IB textbook”

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Melting PointsThe temperature at which a solid becomes a liquid at a

fixed pressure (degrees Kelvin) The temperature at which a crystalline melts depends on the strength of the attractive forces and on the way the particles are packed in the

solid state

• Requires understanding of concepts covered in later topics (this year and next year)

• Know the type of bonding

3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Melting Points

• Don’t worry about the periodic trend!!!• You will need to know the group trends

for Alkali metals and halogens… they’re different!

• Alkali Metals: Melting point decreases down the group– Li (181 oC) to Cs (29 oC)– As the atoms get larger the forces of

attraction between them decrease due to the type of bonding (metallic)

– The “sea of electrons” is further away from the metal ions

ElementElement Melting Melting Point (K)Point (K)

LiLi 453453

NaNa 370370

KK 336336

RbRb 312312

CsCs 301301

FrFr 295295

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Melting Points

• Halogens: Melting point increases down the group– F2 (-220 0C) to I2 (114 oC)– Halogens molecules are held

together with weak van der Waals’ attractive forces due to their non-polar covalent nature

– Larger molecules have more electrons, increasing the strength of the IMF

increases increases

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

ReactivityThe relative capacity of an atom to undergo a chemical

reaction with another atom, molecule or radical.

• Why do atoms react?

• Atoms that are good at becoming stable (getting a full valence shell) are the most reactive

• Don’t worry about the periodic trend because elements on opposite sides of the periodic table can be equally reactive… but for different reasons!!!

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Reactivity of Alkali metalsThe relative capacity of an atom to undergo a chemical

reaction with another atom, molecule or radical.

• How many valence electrons do the alkali metals have?

• Are alkali metals more likely to give up or get electrons?

• What is the name of the property responsible for doing that?

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Reactivity of Alkali metals

• It all has to do with ionization energy since they want to lose electrons!

• Lower ionization energy = more reactive

• Group trend for Alkali metals– Increases as you move down group 1– Since alkali metals are more likely to lose an

electron, the ones with the lowest ionization energy are the most reactive since they require the least amount of energy to lose a valence electron.

• Which alkali metal is the best at losing electrons?• That’s the most reactive alkali metal!

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Reactivity of halogensThe relative capacity of an atom to undergo a chemical

reaction with another atom, molecule or radical.

• How many valence electrons do the halogens have?

• Are halogens more likely to give up or get electrons?

• What is the name of the property responsible for doing that?

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3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization

energies, electronegativities and melting points

Reactivity of halogens

• It all has to do with electronegativity (or electron affinity) since they want to gain electrons!

• Higher electronegativity = lower electron affinity = more reactive

• Group trend for Alkali metals– Decreases as you move down group 17 in the periodic

table – Since halogens are more likely to gain an electron, the

ones with the greatest electronegativity are the most reactive since they are most effective at gaining a valence electron.

• Which alkali metal is best at getting electrons?• That’s the most reactive halogen!

most reactive

least reactive

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IB Topic 3: Periodicity 3.3: Chemical properties

• Discuss the similarities and differences in the chemical properties of elements in the same group.

• Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3.

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3.3.1 Discuss the similarities and differences in the chemical

properties of elements in the same group.

Alkali Metals

React with water &react with many

substances because…

They have the same number of

valence electrons

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3.3.1 Discuss the similarities and differences in the chemical

properties of elements in the same group.

Alkali Metals

2Na(s) + 2H2O(l) 2NaOH (aq) + H2(g)

In the reaction of alkali metals and water, all will:

• move around the surface of the water,

• give off hydrogen gas, • create a basic

solution.

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3.3.1 Discuss the similarities and differences in the chemical

properties of elements in the same group.

Alkali Metals

In the reaction of alkali metals and water, the reactivity will

increase down the group because they get better at getting rid of their valence

electron (the 1st ionization energy

decreases)

So, alkali metals lower down will:

• React more vigorously• React faster

• Give off a flame

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3.3.1 Discuss the similarities and differences in the chemical

properties of elements in the same group.

Alkali Metals

Reaction with halogens

2M(s) + X2 (g) 2MX(s)

where M represents Li,Na,K,Rb, or Cs

Where X represents F,Cl,Br, or I

2Na(s) + Cl2(g) 2NaCl(s)

Reactivity decreases down the group

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3.3.1 Discuss the similarities and differences in the chemical

properties of elements in the same group.Halogens

Halogens are diatomic as gases (two atoms bond together) and called halides when they form ions… These are BrINClHOF

Halogens want to get one electron to fill its outer shell.

Reactivity decreases down the group because electronegativity decreases

Cl2 reacts with Br- and I- Cl2(aq) + 2Br-(aq) 2Cl-(aq) + Br2(l)

Cl2(aq) + 2I-(aq) 2Cl-(aq) + I2(s)

Br2 reacts with I-Br2(aq) + 2I-(aq) 2Br-(aq) + I2(s)

I2 non-reactive with halide ions

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Reactivity of Elements… in action

Alkali Metals: http://www.youtube.com/watch?v=m55kgyApYrY

Halogens:

http://www.youtube.com/watch?v=tk5xwS5bZMA&feature=related

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3.3.2 Discuss the changes in nature, from ionic to covalent and from

basic to acidic, of the oxides across period 3.

Metallic Oxides in Period 3Sodium oxide: Na2O ionic

Magnesium oxide: MgO ionicAluminum oxide: Al2O3 ionic

Metalloid oxide in Period 3Silicon dioxide: SiO2 covalent

Nonmetallic oxides in Period 3Tetraphosphorus decoxide: P4O10 covalent

Sulfur trioxide: SO3 covalent

Dichlorine heptoxide: Cl2O7 covalent

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3.3.2 Discuss the changes in nature, from ionic to covalent and from

basic to acidic, of the oxides across period 3.

Acidic/Basic

Metallic oxides in Period 3 are basicSodium oxide: Na2O + H2O 2 NaOH basicMagnesium oxide: MgO + H2O Mg(OH)2 basicAluminum oxide: Al2O3 + H2O 2 Al(OH)3 amphoteric

Metalloid oxide in Period 3 is acidicSilicon dioxide: SiO2 + H2O H2SiO3 acidic

Nonmetallic oxides in Period 3 are acidicTetraphosphorus decoxide: P4O10 + 6H2O 4H3PO4 acidicSulfur trioxide: SO3 + H2O H2SO4 acidicDichlorine heptoxide: Cl2O7 + H2O 2HClO4 acidicArgon does not form an oxide

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Terms to Know

• Group

• Period

• Alkali metals

• Halogens

• Ionic radius

• Electronegativity

• First ionization energy

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Periodic Table of Video

• http://www.periodicvideos.com/