Periodic Variation in Physical Properties of the Elements H to Ar 38.1The Periodic Table...

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e Elements H to Are Elements H to Ar

38.138.1 The Periodic TableThe Periodic Table

38.238.2 Periodic Variation in Physical Properties Periodic Variation in Physical Properties

of Elementsof Elements

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38.138.1The Periodic TablThe Periodic Tabl

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The Periodic Table The Periodic Table

38.1 The Periodic Table (SB p.2)

• With more and more elements being discovered

needed a way to organize them effectively



• The modern Periodic Table

the basis of the atomic numbers and electronic configurations of eleme

nt


The modern Periodic Table



• The earliest version of the Periodic Table

introduced in 1869

by a Russian chemist called Dimitri Mendeleev


A portion of one of Dimitri Mendeleev’s handwritten drafts of the Periodic Table


Dimitri Mendeleev’s Periodic Table in 1872



• Mendeleev created the first Periodic Table based on atomic masses

• Many elements had similar properties

occurred periodically

the name Periodic Table was used



• The periodic law stated

the chemical and physical properties of the elements vary in a periodic

way with their atomic masses



• Example:

Lithium, sodium, potassium, rubidium and caesium

have similar chemical properties



• Example:

Beryllium, magnesium, calcium, strontium and barium

also have similar chemical properties



• According to Mendeleev’s theory

they could be perfectly arranged by increasing atomic masses

• Some elements did not match perfectly



• Tellurium is heavier than iodine

but the chemical properties of tellurium did not match with th

ose of chlorine and bromine

the chemical properties of iodine did not match with those of sulphur a

nd selenium



• Tellurium should be placed before iodine

even though tellurium was heavier than iodine



• The modern Periodic Table

arranged according to atomic numbers instead of atomic ma

sses



• The modern Periodic Table is divided into

7 horizontal rows called periods

18 vertical columns called groups



• Elements with atoms having the same number of electron shells

put in the same period

• Elements having the same number of outermost shell electrons

put in the same group



• Elements can be classified as

s-block elements

p-block elements

d-block elements

f-block elements


• Group IA and Group IIA elements constitute the s-block

• They are elements with outermost shell electrons occupying the s orbital

1. 1. s s -Block Elements-Block Elements


• Group IA elements have only one outermost shell electron occupying the s orbital

• Examples:

Lithium, sodium, potassium, rubidium, caesium and francium



• They are highly reactive metals

• They are known as the alkali metals



• Group IIA elements have two outermost shell electrons in the s orbital

• Example:

Beryllium, magnesium, calcium, strontium, barium and radium



• They are also chemically reactive

• They are known as the alkaline earth metals



• Elements having electronic configurations from [ ] ns2np1 to [ ] ns2np6

• They include Group IIIA, IVA, VA, VIA, VIIA and 0

2. 2. p p -Block Elements-Block Elements


• Group VIIA elements are all non-metals

• They are known as the halogens



• Group 0 elements are called noble gases

• They have a fully-filled outermost electron shell

gives rise to extra stability

the very stable electronic configuration



• s-Block and p-block elements together are also known as representative elements



• Elements with electronic configurations from [ ] (n – 1)d1ns2 (Group IIIB) to[ ] (n – 1)d10ns2 (Group IIB)

• They are also called transition elements

3. 3. d d -Block Elements-Block Elements


• Two series of f-block elements in which the 4f and 5f orbitals being filled up with 1 to 14 electrons respectively

• They are the lanthanide series and the actinide series

• They are sometimes called inner-transition elements

4. 4. f f -Block Elements-Block Elements


Elements can be classified as s-block elements, p-block elements, d-block elements and f-block elements in the Periodic Table


Check Point 38-1Check Point 38-1

First ionization enthalpy First ionization enthalpy

38.2 Periodic Variation in Physical Properties of Elements (SB p.6)

The first ionization enthalpy of an atom is the energy required to remove one mole of electrons from one mole of its gaseous atoms to form one mole of gaseous ions with one positive charge.



• Energy is required

overcome the attractive forces between the nucleus and the electr

on to be removed

the ionization enthalpy always has a positive value



• The ionization enthalpy of an element

reflects the relative force of attraction between the nucleus and the electron being removed



• Four main factors affecting the magnitude of the ionization enthalpy of an atom:

1. the electronic configuration of an atom;

2. the nuclear charge;

3. the screening effect; and

4. the atomic radius


The first ionization enthalpies of the first 20 elements


Variation in the first ionization enthalpy of the first 20 elements


• The consequence of the increase in nuclear charge with atomic numbers

1.1. General increase in the first ionizatiGeneral increase in the first ionization enthalpy across both Periods 2 aon enthalpy across both Periods 2 and 3nd 3


• At the same time

additional electrons are entering the same electron shell

they have poor screening effect



• In other words

an increase in effective nuclear charge across the periods



• Going across a period

the electrons are drawn closer to the nucleus

more energy is required to remove an electron from the atom

the first ionization enthalpy generally increases across both Periods 2 a

nd 3



• In Period 2

the first ionization enthalpy of boron is lower than that of beryllium

2.2. Irregularities with general increase iIrregularities with general increase in the first ionization enthalpy acrosn the first ionization enthalpy across both Periods 2 and 3s both Periods 2 and 3


• In Period 3

the first ionization enthalpy of aluminium is lower than that of magnesium



• Boron and aluminium have [ ] ns2np1 electronic configurations

easier to remove the outermost p electron



the electron is shielded from the attraction of the nucleus by the completely filled s orbitals (ns2)

The first ionization enthalpies of Group III elements are not very high



• Beryllium and magnesium have a relatively stable electronic configuration

the s orbital is completely filled

a relatively large amount of energy is needed to ionize their atoms



• In Period 2

the first ionization enthalpy of oxygen is lower than that of nitrogen



• In Period 3

the first ionization enthalpy of sulphur is lower than that of phosphorus



• The atoms of oxygen and sulphur have one electron more than the half-filled p sub-shell

when the electronic configuration of half-filled p sub-shell (np3) is attai

ned

extra stability is gained



• A relatively small amount of energy is required to remove the first electron from the atoms of oxygen and sulphur



• The electronic configurations of nitrogen and phosphorus are [ ] ns2np3 (i.e. half-filled p sub-shell)

a relatively stable electronic configuration

more energy is required to remove an electron from their atoms



• The element at the end of each period (i.e. the noble gas)

has a completely filled octet (except helium which has a duplet)

this electronic configuration is very stable

3.3. A sharp drop in the first ionization entA sharp drop in the first ionization enthalpy from one period to the nexthalpy from one period to the next


A large amount of energy is required to remove an electron from their

atoms



• The element at the beginning of the next period (i.e. the Group I element)

has an electron entering a new electron shell

further away from the nucleus



The attractive force between the nucleus and the electron is relatively weak



• This s electron is shielded from the attraction of the nucleus effectively by the inner electron shells

once this electron is removed, a stable electronic configuration is

attained



The first ionization enthalpies of Group I elements are relatively low



• When going down a group

increase in atomic radius

the outermost shell electrons will experience less attraction from the

nucleus

4.4. The first ionization enthalpy decreases The first ionization enthalpy decreases down any group in the Periodic Tabledown any group in the Periodic Table


• There is an increase in the nuclear charge down a group

the outermost shell electrons would experience less attraction from t

he positively charged nucleus

the first ionization enthalpy decreases down a group

4.4. The first ionization enthalpy decreases The first ionization enthalpy decreases down any group in the Periodic Tabledown any group in the Periodic Table

Atomic radius Atomic radius


Atomic radius is used to describe the size of an atom.



• For non-metals

the atomic radii commonly used are the covalent radii



• For metals

the metallic radii are used



Covalent radius is defined as half the internuclear distance between two covalently bonded atoms in a molecule of the element.



Metallic radius is defined as half the internuclear distance between two atoms in a metallic crystal.



• The atomic radius of an atom is governed by two factors:

1. Attraction between the nucleus and the electrons

2. Screening of the outermost shell electrons from the nucleus by inne

r electron shells


• The greater the number of protons in the nucleus

the higher the nuclear charge

1. 1. Attraction between the nuclAttraction between the nucleus and the electronseus and the electrons


• This results in greater attraction between the nucleus and the electrons

the electrons are drawn closer to the nucleus

the atomic radius becomes smaller

1. 1. Attraction between the nuclAttraction between the nucleus and the electronseus and the electrons


• As electrons are negatively charged

repulsion between the outermost shell electrons and the electrons on the inner shells of an atom

the outermost shell electrons are screened from the attraction of t

he nucleus

2. 2. Screening of the outermost shell Screening of the outermost shell electrons from the nucleus by inner elelectrons from the nucleus by inner electron shellsectron shells


• The greater the number of electron shells in the atom

the greater the screening effect



• The outermost shell electrons are less strongly held by the nucleus

the atomic radius becomes larger



The atomic radii of the first 20 elements


Variation in atomic radius of the first 20 elements


• Within a given period

the atomic radii decrease progressively with increasing atomic numbers

an increase in atomic number by one means that one more electron and one more proton are added in the atom



• The additional electron

cause an increase in repulsion between the electrons in the

outermost shell

results in an increase in atomic radius



• The additional proton in the nucleus

cause the electrons to experience greater attractive forces from the

nucleus



• The newly added electron

goes to the outermost shell

is at approximately the same distance from the nucleus

the repulsion between the electrons is relatively ineffective to cause an

increase in atomic radius



• The effect of increasing nuclear charge outweighs the effect of repulsion between the electrons

an increase in effective nuclear charge

the atomic radii of elements decrease across a period



• If we look closer

sharp decrease in atomic radius from the first element to the third eleme

nt of each period

followed by a gradual decrease along subsequent elements



• At the beginning of each period

increasing effective nuclear charge with atomic numbers predomin

ates

greater contraction of the electron cloud



• When more electrons are added to the same electron shell

the effect of repulsion between electrons becomes more significa

nt

the effective nuclear charge increases only slowly towards the

end of the period

the decrease in atomic radius is thus smaller



• Going down a group in the Periodic Table

the atoms have more electron shells occupied

the outermost electron shells become further away from the nucleus



• The outermost shell electrons

more effectively shielded by the inner electron shells from the nuclear charge

decrease in the attractive force between the nucleus and the

outermost shell electrons

the atomic radii of elements increase down a group


Electronegativity Electronegativity


Electronegativity is the relative tendency of an atom to attract bonding electrons towards itself in a covalent bond.


• Pauling assigned electronegativity values to the elements on an arbitrary scale from 0 to 4



• The higher the electronegativity value of an atom

the higher the ability of the atom to attract bonding electrons towa

rds itself in a covalent bond



• Fluorine

the most electronegative element

assigned an electronegativity value of 4.0 in Pauling’s scale



Electronegativity values of the first 20 elements


Variation in electronegativity values of the first 20 elements


• Going across Periods 2 and 3 in the Periodic Table

electronegativity of the elements increases from left to right

the decrease in atomic size



• As the effect of increasing nuclear charge outweighs the screening effect of the electrons in the same electron shell

the bonding electrons are attracted more strongly



• Moving down a group in the Periodic Table

electronegativity of the elements decreases

the increase in atomic size



• With the increase in number of electron shells and greater screening effect

the bonding electrons are attracted less strongly


Melting point Melting point


The melting point of a substance is the temperature at which the substance changes from its solid phase to liquid phase.


• A solid does not melt

unless there is sufficient energy to overcome the forces holding t

he particles together in the solid state



• The amount of energy depends on

1. the magnitude of the attractive forces between the particles

2. how the particles are arranged in the solid



The melting points of the first 20 elements


Variation in melting point of the first 20 elements


• Both lithium and beryllium have a giant metallic structure

• The metallic bond strength increases with the number of outermost shell electrons

1.1. A steady increase in melting point from litA steady increase in melting point from lithium to boron and from sodium to alumihium to boron and from sodium to aluminiumnium


• As lithium has only one outermost shell electron while beryllium has two

the metallic bond strength in beryllium is stronger than that in lithium

beryllium has a higher melting point than lithium



• Boron has a giant covalent structure

• The bonding that holds the boron atoms together is stronger than those of lithium and beryllium

the melting point of boron is higher than those of lithium and berylli

um



• For Period 3 elements

sodium, magnesium and aluminium all have a giant metallic structure



• There is an increase in number of electrons involved in the metallic bond

the strength of metallic bond increases from sodium to aluminium

the melting point increases from sodium to aluminium



• Both carbon and silicon have a giant covalent structure

the atoms are held together by strong covalent bonds

2.2. Carbon and silicon correspond to the maxCarbon and silicon correspond to the maxima in Periods 2 and 3 respectivelyima in Periods 2 and 3 respectively


• A large amount of energy is needed to overcome the strong covalent bonds

the melting points of carbon and silicon are extremely high

2.2. Carbon and silicon correspond to the maxCarbon and silicon correspond to the maxima in Periods 2 and 3 respectivelyima in Periods 2 and 3 respectively


• They all exist as discrete molecules

held together by weak van der Waals’ forces

3.3. The melting points of the elements from The melting points of the elements from nitrogen to neon and from phosphorus tnitrogen to neon and from phosphorus to argon are relatively lowo argon are relatively low


• Only a little amount of energy is needed to overcome the weak van der Waals’ forces

their melting points are relatively low

3.3. The melting points of the elements from The melting points of the elements from nitrogen to neon and from phosphorus tnitrogen to neon and from phosphorus to argon are relatively lowo argon are relatively low



• In Period 3

sulphur exists as S8 molecules in its molecular crystal

phosphorus exists as P4 molecules in its solid molecular crystal



• S8 molecule

a higher molecular mass

a greater surface area for contact with neighbouring molecules



• The van der Waals’ forces between S8 molecules are stronger than those between P4 molecules

the melting point of sulphur is higher than that of phosphorus



• Sulphur

higher melting point than chlorine



• Chlorine

only exists as diatomic molecules

the van der Waals’ forces between S8 molecules are stronger th

an those between Cl2 molecules


Example 38-1Example 38-1 Check Point 38-2Check Point 38-2


Structure and Structure and BondingBonding

A summary of the variations in structure and bonding of elements across both Periods 2


Structure and Structure and BondingBonding

A summary of the variations in structure and bonding of elements across both Periods 3


Structure and BondingStructure and Bonding

• In each of the periods, the structures of the elements changes from

giant metallic structures

followed by giant covalent structures

finally to simple molecular structures


Structure and BondingStructure and Bonding

• Going across the periods from left to right

the bonding of the elements also varies in a repeating pattern

from metallic bonding to covalentbonding

The END

The atomic numbers of tellurium and iodine are 52 and 53 respectively. Why is tellurium heavier than iodine?

AnswerAtomic number of an element is not related to the mass of an atom

of the element. The atomic number of an element is the number of p

rotons in an atom of the element. It is unique for each element. The

mass of an atom of the element is mainly determined by the number

of protons and neutrons in the nucleus. Therefore, tellurium is heavi

er than iodine though the atomic number of tellurium is smaller than

that of iodine.Back


To which block (s-, p-, d- or f-) in the Periodic Table do rubidium, gold, astatine and uranium belong respectively? AnswerRubidium: s-block

Gold: d-block

Astatine: p-block

Uranium: f-block

Back


Which element would have the highest first ionization enthalpy?

AnswerHelium

Back


Which element would have the smallest atomic radius?

AnswerHelium

Back


Why is the melting point of chlorine higher than argon?

AnswerChlorine atom has a higher effective nuclear charge than argon atom,

so the atomic radius of chlorine is smaller than that of argon. Therefore,

the van der Waals’ forces between chlorine molecules are stronger th

an those between argon molecules. Since a higher amount of energy i

s needed to overcome the stronger van der Waals’ forces, the melting

point of chlorine is higher than that of argon.

Back


Considering the trend of atomic radius in the Periodic Table, arrange the elements Si, N and P in the order of increasing atomic radius. Explain your answer briefly.

AnswerIn the Periodic Table, N is above P in Group VA. As the atomic radius increases down a group, the atomic radius of N is smaller than that of P.

Si and P belong to the same period. Since the atomic radius decreases across a period, the atomic radius of P is smaller than that of Si.

Therefore, the atomic radius increases in the order: N < P < Si.

Back


(a) With the help of the Periodic Table only, arrange the elements selenium, sulphur and argon in the order of increasing first ionization enthalpies.Answer

(a) The first ionization enthalpy increases in the order: Se < S < Ar.


(b) Describe and explain the general periodic trend of atomic radius of elements in the Periodic Table.

Answer



(b) Within a given period, the atomic radii decrease progressively with incr

easing atomic numbers. This is because an increase in atomic number

by one means that one more electron and one more proton are added i

n the atom. The additional electron would cause an increase in repulsi

on between the electrons in the outermost shell and results in an incre

ase in atomic radius. The additional proton in the nucleus would cause

the electrons to experience greater attractive forces from the nucleus.

Due to the fact that the newly added electron goes to the outermost sh

ell and is at approximately the same distance from the nucleus, the rep

ulsion between the electrons is relatively ineffective to cause an increa

se in atomic radius. Therefore, the effect of increasing nuclear charge

outweighs the effect of repulsion between the electrons. That means, t

here is an increase in effective nuclear charge. As a result, the atomic r

adii of elements decrease across a period.

(c) With reference to Fig. 38-9 on p.11 (variation in electronegativity value of the first 20 elements), explain why the alkali metals are almost at the bottom of the troughs, whereas the halogens are at the peaks of the plot.

Answer


(c) The alkali metals are almost at the bottom of troughs, indicating that

they have low electronegativity values. It is because their nuclear c

harge is effectively shielded by the fully-filled inner electron shells o

f electrons, and the bonding electrons are attracted less strongly. O

n the other hand, the halogens appear at the peaks. This indicates t

hat they have high electronegativity values. It is because they have

one electron less than the octet electronic configuration. They tend t

o attract an electron to complete the octet, and the bonding electron

s are attracted strongly.

Back


Periodic Variation in Physical Properties of the Elements H to Ar 38.1The Periodic Table...

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Transcript of Periodic Variation in Physical Properties of the Elements H to Ar 38.1The Periodic Table...