Periodic Variation in Physical Properties of the Elements H to Ar

Periodic Variation in Physical Properties of the Elements H to Ar
38 38.1 The Periodic Table 38.2 Periodic Variation in Physical Properties of Elements

38.1 The Periodic Table

The Periodic Table With more and more elements being discovered
38.1 The Periodic Table (SB p.2) The Periodic Table With more and more elements being discovered  needed a way to organize them effectively

The Periodic Table The modern Periodic Table
38.1 The Periodic Table (SB p.2) The Periodic Table The modern Periodic Table  the basis of the atomic numbers and electronic configurations of element

The modern Periodic Table
38.1 The Periodic Table (SB p.2) The modern Periodic Table

The Periodic Table The earliest version of the Periodic Table
38.1 The Periodic Table (SB p.3) The Periodic Table The earliest version of the Periodic Table  introduced in 1869  by a Russian chemist called Dimitri Mendeleev

38.1 The Periodic Table (SB p.3)
A portion of one of Dimitri Mendeleev’s handwritten drafts of the Periodic Table

Dimitri Mendeleev’s Periodic Table in 1872
38.1 The Periodic Table (SB p.3) 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 Table The periodic law stated
38.1 The Periodic Table (SB p.3) The Periodic Table The periodic law stated  the chemical and physical properties of the elements vary in a periodic way with their atomic masses

The Periodic Table Example:
38.1 The Periodic Table (SB p.3) The Periodic Table Example: Lithium, sodium, potassium, rubidium and caesium  have similar chemical properties

The Periodic Table Example:
38.1 The Periodic Table (SB p.3) The Periodic Table Example: Beryllium, magnesium, calcium, strontium and barium  also have similar chemical properties

The Periodic Table According to Mendeleev’s theory
38.1 The Periodic Table (SB p.3) The Periodic Table According to Mendeleev’s theory  they could be perfectly arranged by increasing atomic masses Some elements did not match perfectly

The Periodic Table Tellurium is heavier than iodine
38.1 The Periodic Table (SB p.3) The Periodic Table Tellurium is heavier than iodine  but the chemical properties of tellurium did not match with those of chlorine and bromine  the chemical properties of iodine did not match with those of sulphur and selenium

The Periodic Table Tellurium should be placed before iodine
38.1 The Periodic Table (SB p.3) The Periodic Table Tellurium should be placed before iodine  even though tellurium was heavier than iodine

The Periodic Table The modern Periodic Table
38.1 The Periodic Table (SB p.3) The Periodic Table The modern Periodic Table  arranged according to atomic numbers instead of atomic masses Let's Think 1

The Periodic Table The modern Periodic Table is divided into
38.1 The Periodic Table (SB p.4) The Periodic Table 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

The Periodic Table Elements can be classified as  s-block elements
38.1 The Periodic Table (SB p.4) The Periodic Table Elements can be classified as  s-block elements  p-block elements  d-block elements  f-block elements

1. s -Block Elements Group IA and Group IIA elements constitute the s-block They are elements with outermost shell electrons occupying the s orbital

1. s -Block Elements Group IA elements have only one outermost shell electron occupying the s orbital Examples: Lithium, sodium, potassium, rubidium, caesium and francium

1. s -Block Elements They are highly reactive metals
38.1 The Periodic Table (SB p.4) 1. s -Block Elements They are highly reactive metals They are known as the alkali metals

1. s -Block Elements Group IIA elements have two outermost shell electrons in the s orbital Example: Beryllium, magnesium, calcium, strontium, barium and radium

1. s -Block Elements They are also chemically reactive
38.1 The Periodic Table (SB p.4) 1. s -Block Elements They are also chemically reactive They are known as the alkaline earth metals

2. p -Block Elements Elements having electronic configurations from [ ] ns2np1 to [ ] ns2np6 They include Group IIIA, IVA, VA, VIA, VIIA and 0

2. p -Block Elements Group VIIA elements are all non-metals
38.1 The Periodic Table (SB p.4) 2. p -Block Elements Group VIIA elements are all non-metals They are known as the halogens

2. p -Block Elements Group 0 elements are called noble gases
38.1 The Periodic Table (SB p.4) 2. p -Block Elements 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

2. p -Block Elements s-Block and p-block elements together are also known as representative elements

3. d -Block Elements Elements with electronic configurations from [ ] (n – 1)d1ns2 (Group IIIB) to [ ] (n – 1)d10ns2 (Group IIB) They are also called transition elements

4. f -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

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

Check Point 38-1

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

38.2 Periodic Variation in Physical Properties of Elements (SB p.6) First ionization enthalpy Energy is required  overcome the attractive forces between the nucleus and the electron to be removed  the ionization enthalpy always has a positive value

38.2 Periodic Variation in Physical Properties of Elements (SB p.6) First ionization enthalpy The ionization enthalpy of an element  reflects the relative force of attraction between the nucleus and the electron being removed

38.2 Periodic Variation in Physical Properties of Elements (SB p.6) First ionization enthalpy 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 Let's Think 2

The first ionization enthalpies of the first 20 elements
38.2 Periodic Variation in Physical Properties of Elements (SB p.6) The first ionization enthalpies of the first 20 elements

Variation in the first ionization enthalpy of
38.2 Periodic Variation in Physical Properties of Elements (SB p.7) Variation in the first ionization enthalpy of the first 20 elements

38.2 Periodic Variation in Physical Properties of Elements (SB p.7)
1. General increase in the first ionization enthalpy across both Periods 2 and 3 The consequence of the increase in nuclear charge with atomic numbers

1. General increase in the first ionization enthalpy across both Periods 2 and 3 At the same time  additional electrons are entering the same electron shell  they have poor screening effect

1. General increase in the first ionization enthalpy across both Periods 2 and 3 In other words  an increase in effective nuclear charge across the periods

1. General increase in the first ionization enthalpy across both Periods 2 and 3 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 and 3

2. Irregularities with general increase in the first ionization enthalpy across both Periods 2 and 3 In Period 2  the first ionization enthalpy of boron is lower than that of beryllium

2. Irregularities with general increase in the first ionization enthalpy across both Periods 2 and 3 In Period 3  the first ionization enthalpy of aluminium is lower than that of magnesium

2. Irregularities with general increase in the first ionization enthalpy across both Periods 2 and 3 Boron and aluminium have [ ] ns2np1 electronic configurations  easier to remove the outermost p electron

2. Irregularities with general increase in the first ionization enthalpy across both Periods 2 and 3  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

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

2. Irregularities with general increase in the first ionization enthalpy across both Periods 2 and 3 In Period 2  the first ionization enthalpy of oxygen is lower than that of nitrogen

2. Irregularities with general increase in the first ionization enthalpy across both Periods 2 and 3 In Period 3  the first ionization enthalpy of sulphur is lower than that of phosphorus

2. Irregularities with general increase in the first ionization enthalpy across both Periods 2 and 3 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 attained  extra stability is gained

2. Irregularities with general increase in the first ionization enthalpy across both Periods 2 and 3 A relatively small amount of energy is required to remove the first electron from the atoms of oxygen and sulphur

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

3. A sharp drop in the first ionization enthalpy from one period to the next 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. A sharp drop in the first ionization enthalpy from one period to the next  A large amount of energy is required to remove an electron from their atoms

3. A sharp drop in the first ionization enthalpy from one period to the next 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

3. A sharp drop in the first ionization enthalpy from one period to the next  The attractive force between the nucleus and the electron is relatively weak

3. A sharp drop in the first ionization enthalpy from one period to the next 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

3. A sharp drop in the first ionization enthalpy from one period to the next  The first ionization enthalpies of Group I elements are relatively low

4. The first ionization enthalpy decreases down any group in the Periodic Table When going down a group  increase in atomic radius  the outermost shell electrons will experience less attraction from the nucleus

4. The first ionization enthalpy decreases down 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 the positively charged nucleus  the first ionization enthalpy decreases down a group

Atomic radius Atomic radius is used to describe the size of an atom.
38.2 Periodic Variation in Physical Properties of Elements (SB p.8) Atomic radius Atomic radius is used to describe the size of an atom.

Atomic radius For non-metals
38.2 Periodic Variation in Physical Properties of Elements (SB p.8) Atomic radius For non-metals  the atomic radii commonly used are the covalent radii

Atomic radius For metals  the metallic radii are used
38.2 Periodic Variation in Physical Properties of Elements (SB p.8) Atomic radius For metals  the metallic radii are used

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

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

Atomic radius The atomic radius of an atom is governed by two factors:
38.2 Periodic Variation in Physical Properties of Elements (SB p.8) Atomic radius 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 inner electron shells

1. Attraction between the nucleus and the electrons
38.2 Periodic Variation in Physical Properties of Elements (SB p.8) 1. Attraction between the nucleus and the electrons The greater the number of protons in the nucleus  the higher the nuclear charge

1. Attraction between the nucleus and the electrons
38.2 Periodic Variation in Physical Properties of Elements (SB p.8) 1. Attraction between the nucleus 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

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

2. Screening of the outermost shell electrons from the nucleus by inner electron shells The greater the number of electron shells in the atom  the greater the screening effect

2. Screening of the outermost shell electrons from the nucleus by inner electron shells The outermost shell electrons are less strongly held by the nucleus  the atomic radius becomes larger Let's Think 3

The atomic radii of the first 20 elements
38.2 Periodic Variation in Physical Properties of Elements (SB p.9) The atomic radii of the first 20 elements

Variation in atomic radius of the first 20 elements
38.2 Periodic Variation in Physical Properties of Elements (SB p.9) Variation in atomic radius of the first 20 elements

Atomic radius Within a given period
38.2 Periodic Variation in Physical Properties of Elements (SB p.9) Atomic radius 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

Atomic radius The additional electron
38.2 Periodic Variation in Physical Properties of Elements (SB p.9) Atomic radius The additional electron  cause an increase in repulsion between the electrons in the outermost shell  results in an increase in atomic radius

Atomic radius The additional proton in the nucleus
38.2 Periodic Variation in Physical Properties of Elements (SB p.9) Atomic radius The additional proton in the nucleus  cause the electrons to experience greater attractive forces from the nucleus

Atomic radius The newly added electron  goes to the outermost shell
38.2 Periodic Variation in Physical Properties of Elements (SB p.9) Atomic radius 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

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

Atomic radius If we look closer
38.2 Periodic Variation in Physical Properties of Elements (SB p.10) Atomic radius If we look closer  sharp decrease in atomic radius from the first element to the third element of each period  followed by a gradual decrease along subsequent elements

Atomic radius At the beginning of each period
38.2 Periodic Variation in Physical Properties of Elements (SB p.10) Atomic radius At the beginning of each period  increasing effective nuclear charge with atomic numbers predominates  greater contraction of the electron cloud

Atomic radius When more electrons are added to the same electron shell
38.2 Periodic Variation in Physical Properties of Elements (SB p.10) Atomic radius When more electrons are added to the same electron shell  the effect of repulsion between electrons becomes more significant  the effective nuclear charge increases only slowly towards the end of the period  the decrease in atomic radius is thus smaller

Atomic radius Going down a group in the Periodic Table
38.2 Periodic Variation in Physical Properties of Elements (SB p.10) Atomic radius 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

Atomic radius The outermost shell electrons
38.2 Periodic Variation in Physical Properties of Elements (SB p.10) Atomic radius 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 is the relative tendency of an atom to attract bonding electrons towards itself in a covalent bond.

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

Electronegativity The higher the electronegativity value of an atom
38.2 Periodic Variation in Physical Properties of Elements (SB p.10) Electronegativity The higher the electronegativity value of an atom  the higher the ability of the atom to attract bonding electrons towards itself in a covalent bond

Electronegativity Fluorine  the most electronegative element
38.2 Periodic Variation in Physical Properties of Elements (SB p.10) Electronegativity Fluorine  the most electronegative element  assigned an electronegativity value of 4.0 in Pauling’s scale

Electronegativity values of the first 20 elements
38.2 Periodic Variation in Physical Properties of Elements (SB p.10) Electronegativity values of the first 20 elements

Variation in electronegativity values of the first 20 elements
38.2 Periodic Variation in Physical Properties of Elements (SB p.11) Variation in electronegativity values of the first elements

Electronegativity Going across Periods 2 and 3 in the Periodic Table
38.2 Periodic Variation in Physical Properties of Elements (SB p.11) Electronegativity Going across Periods 2 and 3 in the Periodic Table  electronegativity of the elements increases from left to right  the decrease in atomic size

Electronegativity 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

Electronegativity Moving down a group in the Periodic Table
38.2 Periodic Variation in Physical Properties of Elements (SB p.11) Electronegativity Moving down a group in the Periodic Table  electronegativity of the elements decreases  the increase in atomic size

Electronegativity With the increase in number of electron shells and greater screening effect  the bonding electrons are attracted less strongly

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

Melting point A solid does not melt
38.2 Periodic Variation in Physical Properties of Elements (SB p.11) Melting point A solid does not melt  unless there is sufficient energy to overcome the forces holding the particles together in the solid state

Melting point The amount of energy depends on
38.2 Periodic Variation in Physical Properties of Elements (SB p.11) Melting point 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
38.2 Periodic Variation in Physical Properties of Elements (SB p.11) The melting points of the first 20 elements

Variation in melting point of the first 20 elements
38.2 Periodic Variation in Physical Properties of Elements (SB p.12) Variation in melting point of the first 20 elements

1. A steady increase in melting point from lithium to boron and from sodium to aluminium Both lithium and beryllium have a giant metallic structure The metallic bond strength increases with the number of outermost shell electrons

1. A steady increase in melting point from lithium to boron and from sodium to aluminium 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

1. A steady increase in melting point from lithium to boron and from sodium to aluminium 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 beryllium

1. A steady increase in melting point from lithium to boron and from sodium to aluminium For Period 3 elements  sodium, magnesium and aluminium all have a giant metallic structure

1. A steady increase in melting point from lithium to boron and from sodium to aluminium 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

2. Carbon and silicon correspond to the maxima in Periods 2 and 3 respectively Both carbon and silicon have a giant covalent structure  the atoms are held together by strong covalent bonds

2. Carbon and silicon correspond to the maxima 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

3. The melting points of the elements from nitrogen to neon and from phosphorus to argon are relatively low They all exist as discrete molecules  held together by weak van der Waals’ forces

3. The melting points of the elements from nitrogen to neon and from phosphorus to 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 Let's Think 4

Melting point In Period 3
38.2 Periodic Variation in Physical Properties of Elements (SB p.13) Melting point In Period 3  sulphur exists as S8 molecules in its molecular crystal  phosphorus exists as P4 molecules in its solid molecular crystal

Melting point S8 molecule  a higher molecular mass
38.2 Periodic Variation in Physical Properties of Elements (SB p.13) Melting point S8 molecule  a higher molecular mass  a greater surface area for contact with neighbouring molecules

Melting point 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

Melting point Sulphur  higher melting point than chlorine
38.2 Periodic Variation in Physical Properties of Elements (SB p.13) Melting point Sulphur  higher melting point than chlorine

Melting point Chlorine  only exists as diatomic molecules
38.2 Periodic Variation in Physical Properties of Elements (SB p.13) Melting point Chlorine  only exists as diatomic molecules  the van der Waals’ forces between S8 molecules are stronger than those between Cl2 molecules

Check Point 38-2 Example 38-1

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

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

Structure 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 Bonding Going across the periods from left to right
38.2 Periodic Variation in Physical Properties of Elements (SB p.14) Structure 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 covalent bonding

The END

Let's Think 1 The atomic numbers of tellurium and iodine are 52 and 53 respectively. Why is tellurium heavier than iodine? Answer Atomic 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 protons 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 heavier than iodine though the atomic number of tellurium is smaller than that of iodine. Back

Check Point 38-1 To which block (s-, p-, d- or f-) in the Periodic Table do rubidium, gold, astatine and uranium belong respectively? Answer Rubidium: s-block Gold: d-block Astatine: p-block Uranium: f-block Back

Which element would have the highest first ionization enthalpy?
38.2 Periodic Variation in Physical Properties of Elements (SB p.6) Let's Think 2 Which element would have the highest first ionization enthalpy? Answer Helium Back

Which element would have the smallest atomic radius?
38.2 Periodic Variation in Physical Properties of Elements (SB p.8) Let's Think 3 Which element would have the smallest atomic radius? Answer Helium Back

Let's Think 4 Why is the melting point of chlorine higher than argon? Answer Chlorine 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 than those between argon molecules. Since a higher amount of energy is needed to overcome the stronger van der Waals’ forces, the melting point of chlorine is higher than that of argon. Back

Example 38-1 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. Answer In 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

Check Point 38-2 (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.

Check Point 38-2 (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 increasing atomic numbers. This is because an increase in atomic number by one means that one more electron and one more proton are added in the atom. The additional electron would cause an increase in repulsion between the electrons in the outermost shell and results in an increase 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 shell and is at approximately the same distance from the nucleus, the repulsion between the electrons is relatively ineffective to cause an increase in atomic radius. Therefore, the effect of increasing nuclear charge outweighs the effect of repulsion between the electrons. That means, there is an increase in effective nuclear charge. As a result, the atomic radii of elements decrease across a period.

Check Point 38-2 (c) With reference to Fig 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 charge is effectively shielded by the fully-filled inner electron shells of electrons, and the bonding electrons are attracted less strongly. On the other hand, the halogens appear at the peaks. This indicates that they have high electronegativity values. It is because they have one electron less than the octet electronic configuration. They tend to attract an electron to complete the octet, and the bonding electrons are attracted strongly. Back

Periodic Variation in Physical Properties of the Elements H to Ar

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Periodic Variation in Physical Properties of the Elements H to Ar

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