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Periodic Table

ENGR. YVONNE LIGAYA F. MUSICO

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TOPIC

Definition of Periodic TableHistorical Development of the Periodic

TableThe Periodic Law and Organization of

Elements in a Periodic TablePeriodic Properties and Periodic Trends

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What is Periodic Table?

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Periodic Table

A table of elements written in the order of increasing atomic number and arranged in horizontal rows (periods) and vertical columns (groups) to show similarities in the properties between elements.

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Historical Development of Periodic Table

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Johann Wolfgang Dobereiner (1780 – 1849)

He grouped similar elements with almost equal atomic weights into groups of three or TRIADS

Example:

(Br, I, Cl)

(Sr, Ca, Ba)

(Se, S, Te)

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John A. Newlands(1838 – 1898)

He grouped all elements in the order of their atomic weights.

He then divided the elements into groups of seven elements (noble gases were unknown at that time)

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Lothar Meyer(1830 – 1895)

He plotted a graph showing an attempt to group elements according to atomic weights

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Dmitri Mendeleev(1834 – 1907)

He arranged elements in the order of increasing atomic weights.

In this table, the first two periods had seven elements each. The next periods contained seventeen elements each.

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Dmitri Mendeleev(1834 – 1907)

The discovery of the inert gas during the 1890’s added and additional element to each period.

He left spaces for elements that might someday be discovered.

He studied the properties of the elements and predicted the properties would be discovered elements.

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Henry Moseley(1877 - 1915)

He used X-rays to determine the atomic number (proton) of the elements.

X-rays are light radiation with high frequency and a short wave length.

The higher the atomic number of the elements the shorter the wavelength of X-rays will be when the element is used as a target.

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Henry Moseley(1877 - 1915)

Resolved discrepancies in Mendeleev’s arrangement

He concluded that the elements should be arranged in the order of increasing atomic number

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Periodic Law and Organization of Elements

in a Periodic Table

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Periodic Law

This law states that some of the physical and many chemical properties of the elements are periodic functions of their atomic numbers.

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The Periodic Table

The Modern Periodic TableElements listed in order of increasing atomic

number

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The Periodic Table

Group or Familiesa vertical column of elementscontains elements with similar chemical propertiesthese groups are divided into A and B subgroups.the group number in the A subgroup shows this

number of valence electron

Example:

Phosphorus belong to VA so its valence e- is 5

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The Periodic Table

Periodsa horizontal row of elementsdesignated by numbers 1-7 on the side of the periodic

tableelements are not related chemicallythe period number denotes the number of main energy

level (shell) of the atom.Example:

Sodium is in period 3 so its atom is composed of 3

shells and its outermost shell is the 3rd shell

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The Periodic Table

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The Periodic Table

MetalsMostly solid except Hg which is liquid at room

at temperatureGood electrical and thermal conductivityHigh density, high melting point and boiling

point

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The Periodic Table

MetalsCombine with non-metals to produce saltsDo not combine with each otherTo the left of the stairstep line that in general

separates the metal and non-metals

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The Periodic Table

Metals

Example:

Na, Cu, Fe, Au, Pb

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The Periodic Table

Metals Stairstep line separating metals from non metal

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The Periodic Table

Non MetalsSeveral exist as gases at room temperaturePoor electrical and thermal conductivityLow density, low melting point and boiling

point

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The Periodic Table

Non MetalsCombine with metals to produce saltsSome combine with each otherTo the right of the of the dark stairstep line

that separates the metals from non metals

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The Periodic Table

Non MetalsExample:

C, N, O, S, Cl, He

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The Periodic Table

Non-Metals Stairstep line separating metals from non metal

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The Periodic Table

Metalloidsproperties intermediate between those of

metals and nonmetals

Examples:B, Al, Si, Sb, Sn, Po

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Group or Families in Periodic Table

Alkali metals – Group IAAlkaline Earth Metals – Group IIAHalogens – Group VIIABoron Family – IIIACarbon Family – IVANitrogen Family – VAOxygen Family – VIANoble Gases – VIIIAGroup 0 – very stable configuration with 8 e- in

the outermost shell except He.

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The Periodic Table of the Elements

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Blocks in Periodic Table

Representative Elements

Noble GasesTransition ElementsInner Transition

Elements

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Representative Elements

The A subgroups IA through VIIAThese elements, the outer energy level is

incomplete and the electrons are occupying s or p orbitals.

The electron configuration will be from ns1 to np5 (n is the period number)

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Representative Elements

Examplea) Sodium (Na) at no. = 11, being 3 and & Group IA.

So Na would have one e- in the 3s orbital, configuration in 3s1 with all lower orbitals being completely filled.

b) Nitrogen (N) at no. 7, in the period 2 & in Group VA.

So N would have 2e- in the 2s orbitals and 3e- in the 2p3 with all lower orbital completely filled.

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Noble Gases

Each elements in this group has a completely filled set of s and p orbitals.

The electron configuration for the outermost e- is ns2, np6 (except He, ns2). This is very stable configuration.

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Noble Gases

Example

Krypton (Kr), period 4, group VIIIA

Configuration is 4s2, 4p6 in the outermost energy level.

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Transition Elements

The series having a set of incomplete d orbitals. These elements are the B subgroups. In general, the outermost energy level here will have

an ns2 configuration [except VIB and IB (ns1)]. The outermost electron added to the electrons in the

inner incomplete d orbital corresponds to the group number in the B subgroups (maximum is 8)

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Transition Elements

Examplea) Iron (Fe)

- Period 4, group VIIIB

- Period 4, the last energy level is the 4th

- Group VIIIB, it has an inner incomplete d orbital in the 3rd energy level (this d orbital are of one energy level lower).

b) In case of Chromium (Cr) group VIB, the configuration is 4s1, 3d5.

c) For copper (Cu) group IB, the configuration is 4s1, 3d10

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Inner Transition Elements

Two series of elements from 58 to 71 called the Lathanide series and

belongs to period 6from 90 to 110 called Actinide series series and

belongs to period 7In general, these elements have three

incomplete energy levels since one electron enters an orbital before the set of orbitals begin filling

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Inner Transition Elements

There are many exceptions to the order of filling of orbitals going across. Consider the following examples:

Chlorine (Cl) is a representative element Manganese (Mn) is a transition element Magnesium (Mg) is a representative element Argon (Ar) is a noble gas Uranium (U) is an inner transition element

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Periodic Properties and Periodic Trends

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Periodic Properties and Periodic Trends

Atomic sizeIonization EnergyElectron AffinityElectronegativity

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Atomic Size

The size of the atoms become bigger as the number of shells increases.

Going down any group in the table there is large increase in atomic size.

The increase in the number of energy levels causes the increase in the atomic radius.

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Atomic Size

Going across any period, there is a small but rather decrease in size of the atomic radius.

As the atomic number increases, the nuclear charge becomes greater.

Each electron is attracted towards the nucleus making it closer to the nucleus causing the decrease in the atomic radius

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Atomic Size

1

2

3

4 5

6

7

INCREASES

INCREASES

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Atomic Size

ExampleUsing the periodic table, arrange the following elements in order of decreasing atomic size: Br, Sr, Sn, I, Cs, Ba.

Solution:Going down the group and across the period the order is Cs, Ba, Sr, Sn, I, Br

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Atomic Size

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Ionization Energy

The amount of energy required to remove an electron from an atom.

Going down the group in the table, the ionization energy decreases from one atom to the next.

Going across a period, there is a general increase of ionization energy.

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Ionization Energy

1

2

3

4 5

6

7

INCREASES

INCREASES

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Ionization Energy

Example:Choose the one with the highest ionization energy: Na, Al, Cl, Br

Solution:Cl – since this has the smallest atomic size (that is the valence electron is closer to the nucleus) ionization energy is highest.

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Electron Affinity

The amount of energy released when an electron is added to an atom.

Going down a group of non-metals such as halogen (Group VIIA) the electron affinity decreases.

Going across the period such as from nitrogen to oxygen to flourine, the electron affinity increases.

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Electron Affinity

1

2

3

4 5

6

7

INCREASES

INCREASES

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Electronegativity

It is defined as the tendency of that atom to attract electrons toward itself

Going down a group electronegativity decreases

Going down across the period electronegativity increases

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Electronegativity

Low electronegativity is characteristic of metals The lower electronegativity are found at the lower left

of the Periodic Table. High electronegativity is a characteristic of non-metals. Fluorine is the most electronegative element. Oxygen

is the second. The electronegativity ranges 2.2 to 4.0

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Electronegativity

1

2

3

4 5

6

7

INCREASES

INCREASES

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Electronegativity

Example:Arrange the following elements in order of increasing electronegativity: Ba, Br, I, Sn, Sr

Solution:Looking at position in the periodic table including the group and period, gives the following order.Ba, Sr, Sn, I, BrElectronegativity ranges between 1.8 & 2.1

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Any Question or Comment?