Periodic Table

Larry SchefflerLincoln High School

The Periodic Table-Key QuestionsWhat is the periodic table ?What information is obtained from the table ?How can elemental properties be predicted based on the Periodic Table?

Periodic Table• The development of the periodic table brought a

system of order to what was otherwise an collection of thousands of pieces of information

• The periodic table is a milestone in the development of modern chemistry. It not only brought order to the elements but it also enabled scientists to predict the existence of yet undiscovered

elements.

Early Attempts to Classify Elements

Dobreiner’s Triads (1827)• Classified elements in sets of three

having similar properties.• Found that the properties of the middle

element were approximately an average of the other two elements in the triad.

Dobreiner’s TriadsElement Atomic

MassAverage Density Average

ClBrI

35.579.9126.9

81.21.563.124.95

3.25

CaSrBa

40.187.6137.3

88.71.552.63.5

2.53

Note: In each case, the numerical values for the atomic mass and density of the middle element are close to the averages of the other two elements

Newland’s Octaves -1863John Newland attempted to classify the then 62 known elements of his day.He observed that when classified according to atomic mass, similar properties appeared to repeat for about every eighth elementHis Attempt to correlate the properties of elements with musical scales subjected him to ridicule.In the end his work was acknowledged and he was vindicated with the award of the Davy Medal in 1887 for his work.

Dmitri Mendeleev Dmitri Mendeleev is

credited with creating the modern periodic table of the elements.

He gets the credit because he not only arranged the atoms, but he made predictions based on his arrangement which were shown to be quite accurate.

Mendeleev’s Periodic Table

• Mendeleev organized all of the elements into one comprehensive table.

• Elements were arranged in order of increasing mass.

• Elements with similar properties were placed in the same row.

Mendeleev’s Periodic Table

Mendeleev’s Periodic Table

Mendeleev left some blank spaces in his periodic table. At the time the elements gallium and germanium were not known. He predicted their discovery and estimated their properties

Periodic Table

The Periodic Table has undergone several modifications before it evolved in its present form. The current form is usually attributed to Glenn Seaborg in 1945

Periodic Table Expanded ViewThe Periodic Table can be

arranged by energy sub levels The s-block is Group IA and & IIA, the p-block is Group IIIA - VIIIA. The d-block is the transition metals, and the f-block are the Lanthanides and Actinide metals

The way the periodic table usually shown is a compressed view. The Lanthanides and actinides (F block)are cut out and placed at the bottom of the table.

Periodic Table: Metallic Arrangement

Layout of the Periodic Table: Metals vs. nonmetals

1IA

18VIIIA

12IIA

13IIIA

14IVA

15VA

16VIA

17VIIA

2

33

IIIB4IVB

5VB

6VIB

7VIIB

8 9VIIIB

10 11IB

12IIB

4

5

6

7

MetalsNonmetals

The Three Broad Classes Are Main, Transition, Rare Earth

Main (Representative), Transition metals, lanthanides and actinides (rare earth)

Reading the Periodic Table: Classification

Nonmetals, Metals, Metalloids, Noble gases

Periodic Table: The electron configurations are inherent in the

periodic table

B2p1

1IA

18VIIIA

12IIA

13IIIA

14IVA

15VA

16VIA

17VIIA

2

33

IIIB4IVB

5VB

6VIB

7VIIB

8 9VIIIB

10 11IB

12IIB

4

5

6

7

H1s1

Li2s1

Na3s1

K4s1

Rb5s1

Cs6s1

Fr7s1

Be2s2

Mg3s2

Ca4s2

Sr5s2

Ba6s2

Ra7s2

Sc3d1

Ti3d2

V3d3

Cr4s13d5

Mn3d5

Fe3d6

Co3d7

Ni3d8

Zn3d10

Cu4s13d10

B2p1

C2p2

N2p3

O2p4

F2p5

Ne2p6

He1s2

Al3p1

Ga4p1

In5p1

Tl6p1

Si3p2

Ge4p2

Sn5p2

Pb6p2

P3p3

As4p3

Sb5p3

Bi6p3

S3p4

Se4p4

Te5p4

Po6p4

Cl3p5

Be4p5

I5p5

At6p5

Ar3p6

Kr4p6

Xe5p6

Rn6p6

Y4d1

La5d1

Ac6d1

Cd4d10

Hg5d10

Ag5s14d10

Au6s15d10

Zr4d2

Hf5d2

Rf6d2

Nb4d3

Ta5d3

Db6d3

Mo5s14d5

W6s15d5

Sg7s16d5

Tc4d5

Re5d5

Bh6d5

Ru4d6

Os5d6

Hs6d6

Rh4d7

Ir5d7

Mt6d7

Ni4d8

Ni5d8

Periodic Table Organization------ Groups or Families

Vertical columns in the periodic table are known as groups or families The elements in a group have similar electron configurations

Periodic Table Organization ---- Periods

Horizontal Rows in the periodic table are known as Periods The Elements in a period undergo a gradual change in properties as one proceeds from left to right

Periodic PropertiesElements show gradual changes in certain physical properties as one moves across a period or down a group in the periodic table. These properties repeat after certain intervals. In other words they are PERIODIC

Periodic properties include:

-- Ionization Energy-- Electronegativity-- Electron Affinity-- Atomic Radius-- Ionic Radius

Ionization energy increases across a period because the positive charge increases.Metals lose electrons more easily than nonmetals.Nonmetals lose electrons with difficulty (they like to GAIN electrons).

Ionization energy is the energy required toRemove an electron from an atom

Trends in Ionization Energy

The ionization energy increases UP a group

Because size increases due to an effect known as the Shielding Effect

Trends in Ionization Energy

Ionization Energies

Ionization Energies are Periodic

Electronegativity Electronegativity

is a measure of the ability of an atom in a molecule to attract electrons to itself.

This concept was first proposed by Linus Pauling (1901-1994). He later won the Nobel Prize for his efforts

This concept was first proposed by Linus Pauling (1901-1994). He later won the Nobel Prize for his efforts

Periodic Trends: Electronegativity

In a group: Atoms with fewer energy levels can attract electrons better (less shielding). So, electronegativity increases UP a group of elements.

In a period: More protons, while the energy levels are the same, means atoms can better attract electrons. So, electronegativity increases RIGHT in a period of elements.

Trends in Electronegativity

Electronegativity

Electronegativity

Electron Affinities

Electron Affinities Are Periodic

Electron Affinity v Atomic Number

The Electron Shielding EffectElectrons between the nucleus and the valence electrons repel each other making the atom larger.

The radius increases on going down a group.

Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus.

The radius decreases on going across a period.

The radius increases on going down a group.

Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus.

The radius decreases on going across a period.

Atomic Radius

Atomic RadiusAtomic Radius

The radius decreases across a period owing to increase in the positive charge from the protons. Each added electron feels a greater and greater + charge because the protons are pulling in the same direction, where the electrons are scattered.

Large Small

Atomic Radius

Atomic Radius

Trends in Ion SizesTrends in Ion SizesRadius in pm

Cations

Cations (positive ions) are smaller than their corresponding atoms

Does the size go up or down when gaining an electron to form an anion?

Does the size go up or down when gaining an electron to form an anion?

F,64 pm9e and 9p

F- , 136 pm10 e and 9 p

-

Ion Sizes

CATIONS are SMALLER than the atoms from which they come.The electron/proton attraction has gone UP and so the radius DECREASES.

Li,152 pm3e and 3p

Li+, 78 pm2e and 3 p

+

Ionic RadiusForming a cation.

Ionic Radius for CationsPositve ions or cations are smaller than the corresponding atoms.

Cations like atoms increase as one moves from top to bottom in a group.

Anions

Anions (negative ions) are larger than their corresponding atoms

Ionic Radius-AnionsIonic Radius-Anions

ANIONS are LARGER than the atoms from which they come.The electron/proton attraction has gone DOWN and so size INCREASES.Trends in ion sizes are the same as atom sizes.

Forming an anion.Forming an anion.

F 64 pm9e- and 9p+

F-, 133 pm10 e- and 9 p+

-

Ionic Radii for Anions

Negative ions or anions are larger than the corresponding atoms.

Anions like atoms increase as one moves from top to bottom in a group.

Ionic Radius for an Isoelectronic Group

Isoelectronic ions have the same number of electrons.

The more negative an ion is the larger it is and vice versa.

Summary of Periodic Trends

Properties of the Third Period Oxides

Properties of the Third Period Chlorides

The D Block Elements

The d block elements fall between the s block and the p block.

They share common characteristics since the orbitals of d sublevel of the atom are being filled.

The D Block ElementsThe D block elements include the transition metals. The transition metals are those d block elements with a partially filled d sublevel in one of its oxidation states.

Since the s and d sublevels are very close in energy, the d block elements show certain special characteristics including:

1. Multiple oxidation states

2. The ability to form complex ions

3. Colored compounds

4. Catalytic behavior

5. Magnetic properties

The D Block ElementsThe d electrons are close in energy to the s electrons. D block elements may lose 1 or more d electrons as well as s electrons. Hence they often have multiple oxidation states

Some common D block oxidation states

Multiple Oxidation StatesThere is no sudden sharp increase in ionization energy as one proceed through the d electrons as there would be with the s block. D block elements can lose or share d electrons as well as s electrons, allowing for multiple oxidation states.Most d Block elements have a +2 oxidation State which corresponds to the loss of the two s electrons. This is especially true on the right side of the d block, but less true on the left.

---- For example Sc+2 does not exist, and Ti+2 is unstable, oxidizing in the presence of any water to the +4 state.

Complex IonsThe ions of the d block and the lower p block have unfilled d or p orbitals.

These orbitals can accept electrons either an ion or polar molecule, to form a dative bond. This attraction results in the formation of a complex ion.

A complex ion is made up of two or more ions or polar molecules joined together.

The molecules or ions that surround the metal ion donating the electrons to form the complex ion are called ligands.

Complex IonsCompounds that are formed with complex ions are called coordination compoundsCommon ligands

Complex ions usually have either 4 or 6 ligands.

K3Fe(CN)6 Cu(NH3)42+

Complex IonsThe formation of complex ions stabilizes the oxidations states of the metal ion and they also affect the solubility of the complex ion.

The formation of a

complex ion often has

a major effect on the

color of the solution of

a metal ion.

The D Block Colored Compounds

In an isolated atom all of the d sublevel electrons have the same energy.

When an atom is surrounded by charged ions or polar molecules, the electric field from these ions or molecules has a unequal effect on the energies of the various d orbitals and d electrons.

The colors of the ions and complex ions of d block elements depends on a variety of factors including:

– The particular element– The oxidation state– The kind of ligands bound to the element

Various oxidation states of Nickel (II)

Colors in the D BlockThe presence of a partially filled d sublevels in a transition element results in colored compounds. Elements with completely full or completely empty subshells are colorless, – For example Zinc which has a full d subshell. Its

compounds are whiteA transition metal ion is colored, if it absorbs light in the visible range (400-700

nanometers).If the compound absorbs a

particular wavelength of light its color will be the composite of those wavelengths that it does not absorb.

In other words it shows its complimentary color.

Colors and d Electron TransitionsWhen ligands are attached to transition metal ions, the d orbitals may split into two groups. Some of the orbitals are at a lower energy than the othersThe difference in energy of these orbitals varies slightly with the nature of the ligand or ion surrounding the metal ion

The energy of the transition: ∆E =hn may occur in the visible region. When white light passes through a compound of a transition metal, light of a particular frequency is absorbed as an electron is promoted from a lower energy d orbital to a higher one. The result is a colored compound

Magnetic PropertiesParamagnetism --- Molecules with one or more unpaired electrons are attracted to a magnetic field. The more unpaired electrons in the molecule the stronger the attraction. This type of behavior is calledDiamagnetism --- Substances with no unpaired electrons are weakly repelled by a magnetic field. Transition metal complexes with unpaired electrons exhibit simple paramagnetism. The degree of paramagnetism depends on the number of unpaired electrons

Catalytic BehaviorMany D block elements are catalysts for various chemical reactions

Catalysts speed up the rate of a reaction with out being consumed.

The transition metals form complex ions with ligands that can donate lone pairs of electrons.

This results in close contact between the metal ion and the ligand.

Transition metals also have a wide variety of oxidation states so they gain and lose electrons in oxidation- reduction reactions

Some Common D Block Catalysts

Examples of D block elements that are used as catalysts:

1. Platnium or rhodium in a catalytic converter

2. MnO2 decomposition of hydrogen peroxide

3. V2O5 in the contact process

4. Fe in Haber process 5. Ni in conversion of

alkenes to alkanes

The Periodic Table--Summary

The periodic table is a classification system. Although we are most familiar with the periodic table that Seaborg proposed more than 60 years ago, several alternate designs have been proposed.

Alternate Periodic Tables

Alternate Periodic Tables II