PeriodicTable1

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Periodic Table

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The Periodic Table-Key Questions

What is the periodic table ?

What information does the table provide ?

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Periodic Table

• _{The development of the periodic table brought a}_{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}_{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 elements that had

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Early Attempts to Classify

Elements

Dobreiner’s Triads (1827)

• Classified elements in sets of three

having similar properties.

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Dobreiner’s Triads

Element

Element Atomic Atomic Mass

Mass

(amu)

Average

Average DensityDensity

(g cm

(g cm-3₎₎

Average Average Cl Cl Br Br I I 35.5 35.5 79.9 79.9 126.9 126.9 81.2 81.2 1.56 1.56 3.12 3.12 4.95 4.95 3.25 3.25 Ca Ca Sr Sr Ba Ba 40.1 40.1 87.6 87.6 137.3 137.3 88.7 88.7 1.55 1.55 2.6 2.6 3.5 3.5 2.53 2.53

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Newland’s Octaves -1863

John 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 element

His Attempt to correlate the

properties of elements with musical

scales subjected him to ridicule.

In the end his work was

acknowledged and he was

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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 also made

predictions based on

his arrangements His

predictions were later

shown to be quite

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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}

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Mendeleev’s Periodic Table

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Mendeleev’s Periodic Table

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The Modern Periodic Table

The Periodic Table has undergone several modifications before it evolved in its

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Periodic Table Expanded View

The 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

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Periodic

Table: Metallic

_{Table: Metallic}

Arrangement

1

IA VIIIA18

1 _IIA2 _IIIA13 _IVA14 _VA15 _VIA16 _VIIA17 2

3 _IIIB3 _IVB4 _VB5 _VIB6 _VIIB7 8 _VIIIB9 10 11_IB _IIB12 4 5 6 7

Metals

Nonmetals

Layout of the Periodic Table: Metals vs. nonmetals

.

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The Three Broad Classes are the

Representative, Transition, & Rare

Earth

Main (Representative), Transition metals,

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Additional Groupings in the

Periodic Table

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Periodic Table:

The electron

configurations are inherent in the

periodic table

B B 2p 2p11

1

IA VIIIA18

1 _IIA2 _IIIA13 _IVA14 _VA15 _VIA16 _VIIA17

2

3 _IIIB3 _IVB4 _VB5 _VIB6 _VIIB7 8 _VIIIB9 10 11_IB _IIB12

4 5 6 7 H 1s1 Li 2s1 Na 3s1 K 4s1 Rb 5s1 Cs 6s1 Fr 7s1 Be 2s2 Mg 3s2 Ca 4s2 Sr 5s2 Ba 6s2 Ra 7s2 Sc 3d1 Ti 3d2 V 3d3 Cr 4s1_3d5

Mn 3d5 Fe 3d6 Co 3d7 Ni 3d8 Zn 3d10 Cu

4s1_3d10

B 2p1 C 2p2 N 2p3 O 2p4 F 2p5 Ne 2p6 He 1s2 Al 3p1 Ga 4p1 In 5p1 Tl 6p1 Si 3p2 Ge 4p2 Sn 5p2 Pb 6p2 P 3p3 As 4p3 Sb 5p3 Bi 6p3 S 3p4 Se 4p4 Te 5p4 Po 6p4 Cl 3p5 Be 4p5 I 5p5 At 6p5 Ar 3p6 Kr 4p6 Xe 5p6 Rn 6p6 Y 4d1 La 5d1 Ac 6d1 Cd 4d10 Hg 5d10 Ag

5s1_4d10

Au

6s1_5d10

Zr 4d2 Hf 5d2 Rf 6d2 Nb 4d3 Ta 5d3 Db 6d3 Mo 5s1_4d5

W 6s1_5d5

Sg 7s1_6d5

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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

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Periodic Table Organization

---- Periods

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Periodic Properties

Elements 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:

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Metals lose electrons

more easily than

nonmetals.

Nonmetals lose electrons

with difficulty. (They like

to

GAIN

electrons).

Ionization energy

increases across a period

because the positive

charge increases.

Ionization energy is the energy required to

remove an electron from an atom

Trends in Ionization Energy

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The ionization energy is

highest at the top of a

group. Ionization energy

decreases as the atom

size increases.

This results from an

effect known as the

Shielding Effect

Trends in Ionization Energy

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Ionization Energies of the

Representative Groups

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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.

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Periodic Trends:

Electronegativity

In a group:

In a group: Atoms with fewer Atoms with fewer energy levels can attract

energy levels can attract

electrons better (less shielding). electrons better (less shielding).

So, electronegativity

So, electronegativity increases increases UP

UP a group of elements. a group of elements.

In a period:

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

increases RIGHT

increases RIGHT in a period of in a period of

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Trends in Electronegativity

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Electronegativity

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Electronegativity

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Electron Affinities

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Electron Affinities Are Periodic

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The Electron Shielding Effect

Electrons

between the

nucleus and

the valence

electrons repel

each other

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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

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Atomic 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, whereas the

electrons are scattered.

Large

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Atomic Radius

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Atomic Radius

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Trends in Ion Sizes

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Cations

Cations (positive ions) are smaller

than their corresponding atoms

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CATIONS

are

SMALLER

than the atoms from

which they come.

The electron/proton attraction has gone UP

and so the radius

DECREASES.

Li

0.152 nm 3e and 3p

Li+,

0.078 nm 2e and 3 p

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Ionic Radius for Cations

Positive ions or cations are

smaller than the corresponding atoms.

Cations like

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Anions

Anions (negative ions) are larger

than their corresponding atoms

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Ionic 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.

F 0.064 nm 9e- and 9p+

F-0.133 nm 10 e- and 9 p+

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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?

Ion Sizes

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Ionic Radii for Anions

Negative ions or anions are larger than the

corresponding atoms.

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Ionic Radius for an

Isoelectronic Group

Isoelectronic ions have the same

number of electrons.

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Summary of Periodic Trends

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Properties of the Third Period

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Properties of the Third Period

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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

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The D Block Elements

The 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

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The D Block Elements

The 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

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Multiple Oxidation States

There 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

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Complex Ions

The 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.

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Complex Ions

Compounds that are formed with

complex ions are called

coordination

compounds

Common ligand

s

Complex ions usually have either 4 or 6

ligands.

K

₃

Fe(CN)

₆

Cu(NH

₃

)

₄2+

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Complex Ions

The 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

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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}

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Colors in the D Block

The 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 white

A transition metal ion exhibits color, if it absorbs light in the visible range (400-700

nanometers)

If the compound absorbs a

particular wavelengths of light its particular wavelengths of light its

color is the composite of those color is the composite of those

wavelengths that it does not absorb.wavelengths that it does not absorb. It shows the complimentary color.

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Colors and d Electron

Transitions

The d orbitals may split into two groups so that two orbitals are at a lower energy than the other three

The difference in energy of these orbitals varies slightly with the nature of the ligand or ion surrounding the

metal ion

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.

When the energy of the transition: ∆E =h may occur in

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Magnetic Properties

Paramagnetism --- 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 called

Diamagnetism --- 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

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Catalytic Behavior

Many D block elements are

catalysts for various reactions

Catalysts speed up the rate of a

chemical reaction with out being consumed.

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

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

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Some Common D Block

Catalysts

Examples of D block elements that are

used as catalysts:

1. Platnium or

rhodium is used in a catalytic converter

2. MnO₂ catalyzes the decomposition

of hydrogen peroxide 3. V₂O₅ is a catalyst for

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Alternate Periodic Tables

Although we are most familiar

with the periodic table that

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Alternate Periodic Tables

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Alternate Periodic Tables II

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Alternate Periodic Tables III

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