chap_7-periodic properties of elements

Periodic Properties of elements Chapter 7

Lecture presentationDr. Rajani Srinivasan

Tarleton State University

Contents • Periodic Table – History• Effective Nuclear charge• Sizes of atoms and elements• Ionization Energy• Electron affinities• Properties of Metals, non metals and Metalloids• Periodic trends in Group1 and Group 2 elements• Periodic trends in non metals

History of Periodic table development• Dmitri Mendeleev in Russia and

Lothar Meyer in Germany independently published nearly identical periodic table schemes.

• They arranged the elements in the increasing order of atomic masses.

• Mendeleev pursued this work further and predicted the positions and properties of unknown compound accurately.

• He predicted the properties of Germanium – eka silicon ( below silicon ) and eka -Aluminium -Gallium

Mendeleev’s elements

Modern Periodic table

• After the discovery of Nuclear model of atoms by Rutherford.

• English physicist Henry Mosley gave the concept of “Atomic Number”

• This resulted in existence of modern periodic table. In which elements are arranged in the increasing order of atomic Number

Effective Nuclear chargeThe effective nuclear charge, Zeff, is found this way:

Zeff = Z − S

where Z is the atomic number and S is a screening constant, usually close to the number of inner core electrons.

Other factor that effects the effective nuclear charge is the distance of the electrons from the nucleus

Zeff(Na)= 11-10= +1Since S orbital is nearer to the nucleus so the Zeff= + 2.5

Effective Nuclear chargeThe effective nuclear charge, Zeff, is found this way:

Zeff = Z − S

Slater explained the anomaly in the value of S.

As per Slater’s Rule Electrons having the same value of n as the electron of interest contributes 0.35 of S

Electrons having n-1 as the electron of interest contributes 0.85 of S

Values even smaller than that contributes 1 od S

Trend Zeff

• Effective nuclear charge increases from Left to right across any period

(i.e. from metallic side to non metallic side)• Effective nuclear charge increases down the

column but it is not as significant as across the period.

Sizes of Atoms and Ions• Non bonding atomic radii is

also called vander wall’s radii.

• The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

• This helps in determining the bond length of the covalent bonds between the atoms

Periodic trends

• Atomic Radii increases from top to bottom (along the group)

• Greater the “n” greater the atomic Radii

( number of valence electrons remains the same but shell number increases.)• Along the period the atomic radii decreases

“Number of electrons increases but shell remains the same”

Atomic Radii

Ionic Radii Size of the atoms in the ion depends upon:1) Nuclear Charge 2) Number of electrons 3) Number of orbitals

Cations

• Cations are smaller than their parent atoms:– The outermost

electron is removed and repulsions between electrons are reduced.

Anions

• Anions are larger than their parent atoms”– Electrons are added

and repulsions between electrons are increased.

Periodic trends

• Ions increase in size as you go down a column:– This increase in size is due to the increasing value of n

Cations and anions decreases in size along the period

Isoelectronic Series

In an isoelectronic series, ions have the same number of electrons.Ionic size decreases with an increasing nuclear charge.

IONIZATION ENERGY

• The ionization energy (I) is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion. Unit = KJ/mol– I are Endothermic i.e – Energy values are +ve.

– The first ionization energy(I1) is that energy required to remove the first electron.

– The second ionization energy (I2) is that energy required to remove the second electron, etc.

Ionization Energy

Example:Na(g) Na+ + e-1 I1 = 496 KJ/mol

(Energy associated with this is first ionization energy.

Na + Na + + + e-1 I2 4562 KJ /mol

(Energy associated with this is first ionization energy

Ionization energy• The energy required to remove successive

electron increases• When all the valence electrons are removed

there is a very high increase in the energy to remove an electron.

• Mg(s) Mg+ + e-1 I1= 738KJ

• Mg+ Mg++ + e-1 I2= 1451kJ

• Mg++ Mg+++ + e-1 I3= 7733KJ

Ionization energy

Trends in First Ionization Energies

• As one goes down a column, less energy is required to remove the first electron.– For atoms in the same group, Zeff is essentially the same, but the valence

electrons are farther from the nucleus.

• Generally, as one goes across a row, it gets harder to remove an electron or more energy is required.– As you go from left to right, Zeff increases.

Trends in First Ionization Energies

However, there are two apparent discontinuities in this trend.

Trends in First Ionization Energies

• The first occurs between Groups IIA and IIIA.

from Be [He]2s2 to Boron [He]2s2p1

( Ionization energy decreases then increases)• In this case the electron is

removed from a p orbital rather than an s orbital.– The electron removed is farther

from the nucleus.– There is also a small amount of

repulsion by the s electrons.

Trends in First Ionization Energies• The second discontinuity

occurs between Groups VA and VIA.

From Nitrogen [He]2s22p3 to Oxygen [He]2s22p4

– Ionization energy first increases then decreases

– The electron removed comes from a doubly occupied orbital.

– Repulsion from the other electron in the orbital aids in its removal.

Electronic configuration of cations

Removal of electrons from the atoms follows the following rules

1) They are removed first from the occupied orbitals with largest principal quantum number

Li = 1s22s1 Li + = 1s2

Fe = [Ar]3d64s2 Fe3+ = [Ar] 3d5

Electronic configuration of Anions2) When forming Anions

Electrons are added first to the unoccupied orbitals with lowest principal quantum number or ‘n’ values

F = [He]2s22p5 F- = [He]2s22p6

Electron Affinity

Electron affinity (EA) is the energy change accompanying the addition of an electron to a gaseous atom:

Cl + e− Cl−

EA is exothermic so has negative value for energy changes

Trends in EA

• EA increases as we go from LEFT to RIGHT• Again it has two exceptions – The first occurs between Groups IA and IIA.– The second discontinuity occurs between Groups

IVA and VA.

Electron Affinity

• The added electron must go in a p orbital, not an s orbital.

• The electron is farther from the nucleus and feels repulsion from the s electrons.

Electron Affinity

• Group VA has no empty orbitals.

• The extra electron must go into an already occupied orbital, creating repulsion.

Metals, Non –metals and MetalloidsELEMENTS

Metals Non –metals Metalloids

• Are shiny and lustrous

• Solids are Malleable and Ductile.

• Good conductors of heat and electricity.

• Metallic oxides are ionic and are basic in nature

• Form cations in aqueous solution

• Dull in color• Solids are brittle, some

of them are hard, some of them are soft

• Poor conductors of heat and electricity

• Non metallic oxides are molecular solids and are acidic in nature.

• Forms anions and oxyanions in aqueous solution

Properties between metals and non metals

Example : Si• Looks like a metal

• Brittle like non metal • Intermediate

conductor of heat and electricity

Properties of Metals and non metals

Metals versus Nonmetals• Metals tend to form cations (Low Ionisation energies)• Nonmetals tend to form anions ( Large –ve electronic affinities)

Metals and non metals • Compounds formed between metals and

nonmetals tend to be ionic.

Na + Cl2 2NaCl

• Metal oxides tend to be basic

Metallic oxides react with water forming hydroxides.

M+ O MO

MO + H2O MOH

(

Metallic oxides (soluble or insoluble ) reacts with acids forming Salt and Water NiO(s) + 2HNO3 Ni(NO3)2+ H2O

Non metals

• Non metallic oxides are acidic • When dissolved in water are acidic

C + O2 CO2

CO2 + H2O H2CO3

Water (basic)

Dry ice

Acidic

Group trends

• Alkali metals are soft, metallic solids.

• The name comes from the Arabic word for ashes.

• They are found only in compounds in nature, not in their elemental forms.

• They have low densities and melting points.

• They also have low ionization energies.

Alkali Metals

Reaction with water

Their reactions with water are highly exothermic.

Alkali Metals• Alkali metals (except Li) react with oxygen to form peroxides.• K, Rb, and Cs also form superoxides:

K + O2 KO2

• They produce bright colors when placed in a flame.

Alkaline Earth Metals

• Alkaline earth metals have higher densities and melting points than alkali metals.

• Their ionization energies are low, but not as low as those of alkali metals.

Reactivity with water

• Beryllium does not react with water,

• magnesium reacts only with steam,

• Other alkaline earth metals react readily with water.

• Reactivity tends to increase as you go down the group.

Group 6A

• Also called Chalcogens or “ Ore formers”• Oxygen, sulfur, and selenium are nonmetals.• Tellurium is a metalloid.• The radioactive polonium is a metal.

Sulfur

• Sulfur is a weaker oxidizer than oxygen.

• The most stable allotrope is S8, a ringed molecule.

Group VIIA: Halogens

• The halogens are prototypical nonmetals.• The name comes from the Greek words halos and gennao: “salt formers.”

Group VIIA: Halogens

• They have large, negative electron affinities.– Therefore, they tend to oxidize

other elements easily.

• They react directly with metals to form metal halides.

• Chlorine is added to water supplies to serve as a disinfectant.

Group VIIIA: Noble Gases

• The noble gases have astronomical ionization energies.

• Their electron affinities are positive.– Therefore, they are relatively unreactive.

• They are found as monatomic gases.