Periodic properties of elements

T.M.Sankaranarayanan17.04.2010

Antoine-Laurent de Lavoisier Johann Wolfgang Döbereiner

Born 26 August 1743(1743-08-26)Paris, France

Died8 May 1794 (aged 50)Paris, France

Influences Guillaume-François Rouelle

In 1817, Johann Wolfgang Döbereiner began to formulate one of the earliest attempts to classify the elements. He found that some elements formed groups of three with related properties. He termed these groups "triads".

Alexandre-Emile Béguyer de Chancourtois

(20 January 1820 – 14 November 1886) was a french geologist and mineralogist who was the first to arrange the chemical elements in order of atomic weights, doing so in 1862.De Chancourtois only published his paper, but did not publish his actual graph with the proposed arrangement. Although his publication was significant, it was ignored by chemists as it was written in terms of geology. It was Dmitri mendeleyev's table published in 1869 that became most recognized.

John Alexander Newlands

BornNovember 26, 1837 (1837-11-26)London

Died July 29, 1898 (1898-07-30)

Nationality English

Fields analytical chemistry

Alma mater Royal College of Chemistry

Known for periodic table

john newlands was an English chemist who in 1865 classified the 56 elements that had been discovered at the time into 11 groups which were based on similar physical properties.

Newlands noted that many pairs of similar elements existed which differed by some multiple of eight in atomic weight. However, his law of octaves, likening this periodicity of eights to the musical scale, was ridiculed by his contemporaries. It was not until the following century, with Gilbert N. lewis’ valence bond theory (1916) and irving langmuir’s octet theory of chemical bonding[ (1919) that the importance of the periodicity of eight would be accepted

Dmitri Mendeleev in 1897

Born8 February 1834(1834-02-08)Verhnie Aremzyani, Russian Empire

Died

2 February 1907 (aged 72)St. Petersburg, Russian Empire

Nationality Russian

Fields Chemistry, physics and adjacent fields

Alma mater Saint Petersburg University

Notable students

Dmitri Konovalov, Gemilian, Valery, Baykov, Alexander

Known for Inventing the Periodic table of chemical elements

MENDELEEF’S PERIODIC TABLEMendeleef proposed following periodic law and arranged elements in the increasing order of atomic masses.

Mendeleef’s Periodic law

"Physical and chemical properties of elements are periodic functions of their atomic weights"

The elements are arranged in horizontal rows and vertical columns.

•The horizontal rows are called periods. The number of elements in a period may vary. The first three periods of the Mendeleef table are called as short periods. The other periods are known as long periods.

•The vertical columns in the table are known as groups or families. The groups are sub divided into two subgroups ‘A’ and ‘B’. The elements arranged in a group possess similar properties.

•The long period of the Mendeleef periodic table consists of two rows of elements called as series.

* He considered the similarities in the formulae and the properties of the compounds formed by the elements.

Mendeleef observed that elements with similar properties had

(i) Either almost the same atomic weights.e.g. Fe(56), Co(59), Ni(59)Os(191), Ir(193), Pt(195)(ii) Or atomic weights which showed a constant increasee.g. K(39), Rb(85), Cs(133)Ca(40), Sr(88), Ba(137)

The elements with low atomic weights are called typical elements. These are arranged in three short periods of the periodic table.

•Group VIII of the Mendeleef table contains three triads, namely, (Fe, Co, Ni and Ru, Rh, Pd and Os, Ir, Pt). These triads are called transition elements which include Sc(21) to Zn (30), lanthanides and actinides.

* From a study of adjacent elements and their compounds, Mendeleef was able to predict the characteristics of certain elements which were found to be very accurate. e.g.1) Eka Al - Gallium 2) Eka Si - Germanium 3) Eka B - Scandium

Comparison of properties predicted by Mendeleef with those observed.

Periodic Table

He corrected the atomic weights of some elements like Beryllium, Indium, Uranium.

Limitations :Zero group elements were not known at the time of Mendeleef.•In every group lanthanides were placed though they have dissimilar properties.

* At some places it violates the increasing order of atomic weight rule. Ar40 & K39 Co59 & Ni58 Te128 & I127 Th 232 & Pa 231These four pairs are called anomalous pairs or inverted pairs.

* Elements with dissimilar properties were grouped together.E.g. ' Th' is placed in III group and Ag is placed in I group.

* Atomic mass is taken as fundamental property.

MODERN LONG FORM OF PERIODIC TABLEModern periodic table was constructed by Neils Bohr based on modern periodic law.

Modern periodic lawThe chemical and physical properties of elements are the periodic functions of their atomic numbers and electronic configurations.

Above law was proposed by Moseley. He found the relation between atomic numbers (z) and the frequencies () of x- rays produced by them.

Where a & b are constants characteristic of elements

Periodic Table Summary

•The subdivisions of the periodic table are periods, groups, and classes. The horizontal rows are called periods. The vertical columns are called groups.•The entire table consists of three classes: metals, non-metals, and semimetals.The subdivisions of the periodic chart have been explained such that the student should be able to identify them if given a periodic table.•Elements of the same group share certain physical and chemical characteristics. Examples of the characteristics of several groups are listed below.•Group 0 contains elements that are unreactive gases.•Group IA contains elements that are chemically active soft metals.•Group VIIA contains elements that are chemically active nonmetals.•Groups IB through VIIB and VIII are called transition groups and•elements found in them display properties of metals.•The valence of an atom is defined as the number of electrons an element•gains or loses, or the number of pairs of electrons it shares when it interacts with other elements.

Number of elements in each period

Development of the Periodic Table

• Mendeleev: The First Periodic Table– Arranged elements by increasing atomic

mass– Noticed a regular pattern in chemical and

physical properties– Blank spaces – Predicted the existence of

elements not yet discovered based on properties

• Moseley– Determined atomic number of elements by

determining nuclear charge– Arranged table by increasing atomic

number – remains this way today

The Modern Periodic Table

• Periods – Horizontal Rows • Groups/Families – Vertical Columns• Elements in groups have similar

physical and chemical properties• Periodic Law: When elements are

arranged in order of increasing atomic number, there is a regular pattern in their chemical and physical properties

Electron Configurations & Periodicity

• Electrons – play a large part in determining physical & chemical properties of elements

• Arrangements related to position on the table– Groups have the same outer orbital

configuration in consecutively higher energy levels

– Periods have outer electrons in the same energy level

Classifications• Noble Gases: elements in which outermost s

and p sublevels are filled [full octet of valence e-]

• Representative (Main Group 1A-7A) elements: outermost s or p sublevels are only partially filled. The number of valence electrons can be determined by the group number

• Transition Metals – elements in which s and nearby d sublevels contain electrons

• Inner Transition Metals – elements in which s and nearby f sublevels contain electrons

Reading Configurations from the Periodic Table

• Read the table like sentences in a paragraph from left to right & top to bottom until a particular element of interest is reached.

• Each period corresponds to a principle energy level, each group corresponds to a particular sublevel and orbital position.– Subtract 1 from the principle energy level

for transition metals (d-block)– Subtract 2 from the principle energy level

for inner transition metals (f-block)

Trends in Atomic Size• Atomic Radius: half of the distance

between the nuclei of two atoms of the same element bonded together.

• Group Trends: Generally increase as one goes down a group in the periodic table– Electrons occupy consecutively higher energy

levels, farther from the nucleus– Nuclear charge increases, but outermost orbitals

are larger due to distance– Shielding effect increases with additional

number of occupied orbitals between the outermost electrons and the nucleus

Atomic Radius Cont.

• Period Trends: Decreases left to right across the periodic table

• Electrons are added to the same principle energy level

• Increased nuclear charge “pulls” them closer to the nucleus

Trends in Atomic Radius

• Atomic Size increases from top to bottom within a group.

• This is due to shielding effect. The inner electrons shield the outer electrons from the nucleus. So, the nucleus has less pull. Therefore, the atom gets bigger down a group.

Trends in Atomic Radius• Atomic size usually

decreases across a period from left to right.

• This is due to the nuclear charge (which is positive) pulling on the electrons (which are negative.)

• The nuclear charge increases causing it to have more pull. It pulls the energy levels in making the atom smaller.

Trends in Ionic Radius• An Ion is an atom or a group of

atoms with a charge.• Positive ions form when

electrons are lost. (cations)• Negative ions form when

electrons are gained. (anions)• Cations are always smaller

than the atoms from which they form.

• Anions are always larger than the atoms form which they form.

Sodium Ion Chlorine Ion

Trends in Ionic Radius

• In general, ionic radius increases down a group due to the electron shielding.

0

20

40

60

80

100

120

140

160

180

Group Trend

Lithium

Sodium

Potassium

Rubidium

Cesium

Trends in Ionic Radius

• Ionic radius decreases across a period for the cations, and decreases across a period for the anions due to nuclear charge.

0

20

40

60

80

100

120

140

160

Period Trend

Lithium

BerylliumBoron

Carbon

Nitrogen

Oxygen

Fluorine

Trends in Ionic Size• Positive ions are smaller than the neutral atoms

from which they are formed– Electrons are lost from energy levels farthest

from the nucleus– Remaining electrons pulled closer to the

nucleus• Negative ions are larger than the neutral atoms

from which they are formed– Electrons are gained, resulting in smaller

effective nuclear charge for the greater number of electrons

– Repulsive forces between electrons increases across the periods

Ionic Size Trends Cont.

• The trends follow the same pattern as atomic size–Larger from top to bottom–Smaller from left to right

Trends in Ionization Energy

• Ionization Energy: energy needed to remove an electron from a gaseous atom

• Group 1: easily loses its 1 valence electron– Low first ionization energy– Second ionization energy will be very

high since it is “happy” with losing 1 electron.

• Group 2: easily lose 2 valence electrons– Low first and second ionization

energies– High third ionization energy

Trends in Ionization Energy

• The energy required to remove an electron from an atom is called ionization energy.

• First Ionization energies are the energies required to remove the first electron.

• Second Ionization energy is the energy required to remove the second electron, and so on…

Trends in Ionization Energy

Atomic Number

Symbol Name 1st 2nd 3rd 4th

1 H Hydrogen 1312

2 He Helium 2372 5250

3 Li Lithium 520 7298 11815

4 Be Beryllium 899 1757 14848 21006

5 B Boron 800 2427 3659 25025

• Lithium only has 1 electron to lose in order to bond. Notice that between 1st and 2nd ionization energies that the number goes up. Because once lithium has lost it’s one electron it becomes stable. Elements like being stable. So, the ionization energy increases.

Trends in Ionization Energy

• First Ionization energies tend to decrease from top to bottom within a group due to electron shielding.

Trends in Ionization Energy• First Ionization energies tend to increase from

left to right across a period due to increased nuclear charge.

Ionization Energy Trends

• Group Trends: decreases from top to bottom on the periodic table– Outermost electrons are farther from the

effect of the nuclear charge and therefore easier to remove

– Shielding effect increases down the table• Period Trends: increases from left to right

– Nuclear charge is increasing with no increase in shielding effect

– Outermost electrons are closer to the nucleus

Electron Affinity Trends

• Electron Affinity: energy change that accompanies the addition of an electron to a gaseous atom– Most electron affinities are negative

because most elements release energy when they become negative ions

• Group: Decreases from top to bottom• Period: Increases left to right

Electron Affinity Trends(Same as for Ionization Energy)

• Group Trends: decreases from top to bottom on the periodic table• Outermost electrons are farther from the

effect of the nuclear charge and therefore easier to remove

• Shielding effect increases down the table• Period Trends: increases from left to right

• Nuclear charge is increasing with no increase in shielding effect

• Outermost electrons are closer to the nucleus

Electronegativity

• Electronegativity: tendency for an atom to attract electrons to itself when it is chemically combined with another element

• Group: decreases top to bottom• Period: increases left to right

Electronegativity Trends

(Similar to Ionization Energy)• Group Trends: decreases from top to

bottom on the periodic table• Outermost electrons are farther from the

effect of the nuclear charge and therefore more difficult to attract

• Shielding effect increases down the table• Period Trends: increases from left to right

• Nuclear charge is increasing with no increase in shielding effect

• Outermost electrons are closer to the nucleus and attracted more easily

Trends in Electronegativity

• Is the ability of an atom of an element to attract electrons when the atom is in a compound.

• Notice that noble gases have no values. This is due to the fact that most do not form compounds.

Li

1.0

Be

1.5

B

2.0

C

2.5

N

3.0

O

3.5

F

4.0

Na

0.9

Mg

1.2

Al

1.5

Si

1.8

P

2.1

S

2.5

Cl

3.0

K

0.8

Ca

1.0

Ga

1.6

Ge

1.8

As

2.0

Se

2.4

Br

2.8

Rb

0.8

Sr

1.0

In

1.7

Sn

1.8

Sb

1.9

Te

2.1

I

2.5

Electronegativity Values

Trends in Electronegativity

• In general, electronegativity values decrease from top to bottom within a group due to the shielding effect. The inner electrons shield the nucleus from being able to attract electrons.

Trends in Electronegativity

• In general, electronegativity values tend to increase from left to right across a period due to an increase in nuclear charge.

General Explanations

• Shielding effect plays a major role in how strongly the nucleus can pull on its outermost electrons – This mainly effects group trends.

• Nuclear charge is a measure of the strength of a nucleus’ pull - This greatly effects period trends because shielding effect is not an issue across a period.

Effective Nuclear charge• The effective nuclear charge, is the

net positive charge experienced by an electron in a multi-electron atom. The term "effective" is used because the shielding effect of negative electrons prevents higher orbital electrons from experiencing the full nuclear charge. It is possible to determine the strength of the nuclear charge by looking at the oxidation number of the atom.

Summary of Trends

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