Unit 2 Electronic

8/6/2019 Unit 2 Electronic

1/22

Unit 2-1

Unit 2 The Electronic Structure of Atoms and

the Periodic Table

Section 2.1 Atomic Emission Spectrum

( 1 ) Energy associated with light

Light is a form of energy. It has dual properties : wave property and particle property.

1.Wave property

Light is an electromagnetic wave. It travels with velocity c ( c = 3.00 x 108

m s-1

). The frequency

and the wavelength of the wave are related by :

That which we call visible light is an electromagnetic wave of wavelength 400 to 700 nm

(1 nm = 10-9 m ). It is a very narrow component of a large group of radiation types comprising theelectromagnetic spectrum which includes radio waves, microwaves, infra-red, visible light, ultraviolet,

X-rays, gamma rays and cosmic rays.

2. Particle property

Light consists of discrete particles. Each discrete particle is an energy packet called photon.

The energy of a photon depends on the frequency of light wave and is gives by the famous Plancks

equation :

The proportionality constant h is called the Plancks constant which is equal to 6.63 x 10-34

J s.

The greater the frequency the light, the shorter the wavelength, and the greater the energy of a photon.

Example: Calculate the energy of 1 mole of photons with wavelength 650 nm.


2/22

Unit 2-2

( 2 ) The uniqueness of atomic emission spectra

Some gaseous materials emit light when they are subject to large potential differences in electric

discharge tubes. Neon advertising signs work in this way and the yellow sodium street lamps are

discharge tubes containing sodium vapour.

If the light emitted by these substances is examined using a spectroscope, it is found not to consist of

a continuous range of colours like the spectrum of white light or the colours in a rainbow. Instead, thelight emitted by these substances is composed of separate lines of different colour. This kind of

spectrum is called a line emission spectrum.

The atoms of each element have a unique arrangement of electrons with definite energy levels.

When the atoms are excited, electrons are brought to higher energy levels. As the electrons return to

lower energy levels, the atom emits radiations of definite wavelengths. This results in a unique emission

spectrum which can provide useful information about the atom.

The atomic emission spectra of sodium and hydrogen shown below illustrate the unique nature of thespectra. Each element has its own pattern of lines in its emission spectrum.


3/22

Unit 2-3

( 3 ) Characteristics of the atomic emission spectrum of hydrogen

When a sample of hydrogen gas is heated to a high temperature or when an electric discharge is

passed through it at low pressure, the hydrogen molecules absorb energy and split into atoms. The

excited hydrogen atoms contain excess energy and are unstable. Their electrons are promoted from the

ground state to a higher energy level. When the excited electrons lose energy and fall back to the ground

state, light of various wavelengths will be emitted to produce the atomic emission spectrum of hydrogen.

The atomic emission spectrum of hydrogen consists of coloured bright lines in dark background.

The spectrum consists of several series ofdiscrete lines which converge in different parts of the

electromagnetic spectrum :

Lyman series in the ultraviolet region,

Balmer series in the visible region,

Paschen series in the infrared region, and

Brackett and Pfund series in the far infrared region.

Within each series, the lines get closer together and converge towards high frequency end of thespectrum. Finally, the lines merge into a continuum of light (light of the highest frequency). This is

the convergence limits. The following figure shows the Balmer series of hydrogen :


4/22

Unit 2-4

( 4 ) Interpretation of the atomic emission spectrum of hydrogen

1. Bohrs model of hydrogen atom

In 1913, Neils Bohr suggested the model of hydrogen atom with reference to quantum theory :

The lowest energy state, when the hydrogen atom has least energy, is called the ground state.

Hydrogen atoms with more energy than they possess in ground state are said to be excited. When a

hydrogen atom is excited, the electron temporarily occupies an energy level further from the nucleus.

As an electron falls back to a lower energy level, a discrete amount of energy is emitted as radiation.

This produces a line in the atomic spectrum.

Energy level is quantised according to quantum theory.

The electrons can only exit in certain fixed energy levels. It cannot possess energy of intermediate

magnitudes. When hydrogen atoms absorb or emit energy, electrons move from one energy level to

another. Only transitions from one level to another are possible. Transitions of intermediate energy are

not allowed. The energy of an electron can only be lost or gained in small packets called quanta. The

amount of energy involved is exactly equal to a photon (a quantum of energy = h).

2. Plancks relationship

The energy of transition can be calculated from the difference in energy between two energy levels.

If the frequency of light emitted is , the difference between two energy levels is related by Plancks

equation :

Where Einitial Efinal ( E ) = quantum of energy

= frequency of light (Hz, s-1)

h = 6.63 x 10-34

J s (Plancks constant)

c = 3.00 x 108

m s-1

(velocity of light)

= wavelength of light (m)

E is proportional to the frequency of radiation and is inversely proportional to wavelength. When

an electron falls from a higher energy level to a lower energy level, it emits a quantum of energy E inradiation ofdefinite frequency, . Since E for the change is always the same in a given atom, must

be a constant. Therefore, radiation always has the same energy and is always of the same frequency for

this particular electron transition. The atomic emission spectrum consists ofdiscrete lines, i.e.

discontinuous.


5/22

Unit 2-5

3. Balmer series

The following figure shows the energy levels in a hydrogen atom and the electron transitions which

produce the lines in the visible region of the atomic emission spectrum of hydrogen :

Note that :

a. The electronic energy levels are numbered (n = 1, n = 2, n = 3, etc.). These numbers are sometimes

referred to as the principal quantum numbers for the energy levels. The level of lowest energy is

given the principal quantum number 1, the next lowest 2, and so on.

b. The coloured lines in the Balmer series are caused by electron transitions from the higher level to thelevel n = 2. For example, transitions from n = 3 to n = 2 result in a red line (at frequency 4.568 x 1014

Hz) in the hydrogen spectrum, while transitions from n = 4 to n = 2 produce a blue line.

Question : Calculate the difference in energy between energy levels of n = 3 and n = 2.

c. As the energy levels get closer and eventually come together, it follows that the spectral lines also getcloser and eventually come together. This particular frequency is called the convergence limit.


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Unit 2-6

4. Lyman Series, Paschen series, Brackett series and Pfund series

When transition occur from the higher levels to the lowest energy level ( n = 1 ), more energy is

released than with transitions to the n = 2 level. Consequently, these lines appear at higher frequencies

in the spectrum. In the case of hydrogen, transitions to the n = 1 level result in lines in the ultraviolet

region of the spectrum and is known as the Lyman series. Similarly, electron falls back to the third,

fourth, and fifth energy levels gives the Paschen, Brackett, and Pfund series respectively.

The wavelengths and frequencies of the discrete lines of the various series was found by Rydberg

experimentally to obey the following relationship :

where RH = 1.0968 x 107

m-1

.

Spectral

series

Final energy

level (n1)

Initial energy

level (n2)

Lyman

Balmer

Paschen

Brackett

Pfund

5. Intensity of the spectral lines

The spectral lines in the atomic emission spectrum of hydrogen are not of equal intensity because the

intensity of a particular spectral line depends on the number of electrons making the electronic transition

corresponding to a particular frequency. It is unlikely that the number of electrons making various

electronic transitions is the same.

6. Convergence limit and ionisation energy

If sufficient energy is given to an atom, it is possible to excite an electron just beyond the highest

energy level. In this case the electron will escape and the atom becomes as ion. Ionisation has taken

place. By determining the frequency at which the converging spectral line come together, the ionisationenergy of an atom can be found.

Question : An accurate value for the frequency at the convergence limit of Lyman series for hydrogen

is 3.27 x 1015

Hz. Calculate the ionisation energy of hydrogen.


7/22

Unit 2-7

Section 2.2 Atomic Orbital

( 1 ) The wave nature of electrons

In 1924, a French Physicist Louis de Broglie postulated that matter could have both particle and

wave properties. For the case of an electron, this postulate was later confirmed by the experiments

performed by C. Davisson with his partner L. Germer, and by G. P. Thomson. When they passed a beam

of electrons through a crystal and through a gold foil, they obtained a diffraction pattern on a screensimilar to that observed when X-ray was used. This piece of evidence demonstrated that electrons did

indeed possess wave properties. Following this, studies in quantum mechanics have shown that

electrons are not localized in fixed orbits, and as a matter of fact, we can find only describe the location of

an electron in terms of probability of finding it in a certain position at any time.

The probability, considered over a period of time, gives an average picture of how the electron

behaves. The picture does not show the actual location of the electron at any given time. Instead, it

shows where the electron is most likely to be at any time. The total picture of the probability of finding

an electron at various points in spaces is called an orbital. The following figure shows the probability

distribution of the electron in a hydrogen atom :

Note that the denser the dots, the more likely the electron will be in that region. Nowhere is the

probability equal to zero. Even at points at very great distances from the nucleus there is some

probability, although it is small.

Atomic orbital

The volume of the space in which there is a 95% chance of finding the electron is called the atomic

orbital. There is a 5% probability that the electron will be outside this volume of space at a giveninstant. An atomic orbital can be viewed as a representation of a region within which there is a high

probability of finding an electron.

Figure below depicts the boundary surface for the electron distribution in hydrogen atom :


8/22

Unit 2-8

( 2 ) Electron shells, subshells and number of orbitals

In quantum mechanics, the electron is treated as wave. A wave function is used to describe its

motion. Solutions of the wave equation for electrons show electronic energy levels are quantized and

described by quantum numbers. The most important quantum is the principalquantumnumbern

which determines the energy and distance of electron from the nucleus.

The term electron shell is used for a group of orbitals with the same principal quantum number. Asubshell is a group of orbitals with the same energy.

For n = 1, electrons are in the lowest energy level ( Kshell ) and are nearest to the nucleus. There

is one sub-level (subshell) and one kind of orbital : the 1sorbital.

For n = 2, electrons are in the second energy level ( Lshell ). There are two sub-levels (subshells)

and two kinds of orbitals : the 2s and 2porbitals.

For n = 3, electrons are in the third energy level ( Mshell ). There are three sub-levels (subshells)

and three kinds of orbitals : the 3s, 3p and 3dorbitals.

For n = 4, electrons are in the fourth energy level ( N shell ). There are four sub-levels (subshells)

and four kinds of orbitals : the 4s, 4p , 4d and 4forbitals.

principal quantum

number n

electron

shell subshell

number of

orbitals name of orbitals

In general, the total number of orbitals for each value of principal quantum number n is n2

.

Each orbital can be occupied by two electrons with opposite spins at most.

n total number of orbitals maximum number

of electrons


9/22

Unit 2-9

( 3 ) Shapes of s and p orbitals

s orbital

The shape of an s orbital is spherically symmetrical about the nucleus. The orbital has no preferred

direction. The probability of finding an electron at a distance r from the nucleus is the same in all

directions. The following figures show the shapes of a 1s orbital and a 2s orbital.

p orbitals

There are three p orbitals for each electron shell except n = 1. They are all dumb-bell in shape. p

orbitals are directional and situated along three coordinate axes (x, y, and z). The px, py and pz orbitals

are perpendicular to one another. They have the same energy and are collectively referred to as the p

subshell. There is zero probability of finding a p electron at the nucleus. Figures below depict three 2p

orbitals : 2px, 2py and 2pz.

d orbitals

There are five d orbitals for each d subshell. Three of them has the shapes shown below :


10/22

Unit 2-10

( 4 ) The relative energies of orbitals

One-electron system

For a one-electron atom, such as hydrogen atom or one-electron ions, such as He+

, Li2+

, the relative

energy level diagram of orbitals is shown in the following diagram. Notice that all subshells of the same

energy level have the same energy.

Many-electron system

For a many-electron atom, the energies of subshells increase in the order s< p< d< f . The relative

energy level diagram of orbitals is shown in figure :


11/22

Unit 2-11

Section 2.3 Electronic Configuration

( 1 ) Building up principles

The distribution of electrons in an atom is called electronic configuration. The electronic

configuration can be expressed in two ways :

1. notations using 1s, 2s, 2p, etc.;2. electrons-in-boxes diagram, which tells how electrons occupy different orbitals with spins shown.

In electron-in-boxes diagrams,

represents an orbital, which can accommodate one electron or 2 electrons at most.

represents one electron which occupies an orbital singly.

represents two electrons having opposite spins.

Building up of electronic configurations is based on three principles.

1. Aufbau principle : electrons enter the orbitals in order of ascending energy.

1s


12/22

Unit 2-12

( 2 ) Electronic configurations of atoms from H to KrThe building up of electronic configurations follows the three principles. In addition, extra

stability is gained by electronic configurations in which orbitals of the same energy are either exactly

half-filled or full-filled.

Exactly half-filled subshells (e.g. np3, nd

5) attain extra stability because electrons are evenly

distributed in different orbitals. This minimizes electronic repulsion.

Full-filled subshells (e.g. ns2, np

6, nd

10) attain extra stability because extra energy is needed to break

the spin paired arrangement in a full-filled subshell.

Hydrogen H (Z = 1) 1s1

(1)

1s 2s 2p

Helium He (Z = 2) 1s2

(2)

1s 2s 2p

The second period elements :

Lithium Li (Z = 3) 1s22s

1(2.1)

1s 2s 2p

Beryllium Be (Z = 4) 1s22s

2(2.2)

1s 2s 2p

Boron B (Z = 5) 1s22s

22p

1(2.3)

1s 2s 2p

Carbon C (Z = 6) 1s22s

22p

2(2.4)

1s 2s 2p

Nitrogen N (Z = 7) 1s22s22p3 (2.5)1s 2s 2p

Note that : According to Hunds rule, the 7th electron occupies another 2p orbital and the 5th, 6th and 7th

electrons have parallel spins. The 2p subshell is now exactly half-filled.

Oxygen O (Z = 8) 1s22s

22p

4(2.6)

1s 2s 2p

Fluorine F (Z = 9) 1s22s

22p

5(2.7)

1s 2s 2p


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Unit 2-13

Neon Ne (Z = 10) 1s22s

22p

6(2.8)

1s 2s 2p

Note that : The 2s and 2p subshells of neon is completely filled. This electronic configuration is very

stable. ( ns2np

6is known as the stable noble gas electronic configuration)

The third period elements :

Sodium Na (Z = 11) 1s22s

22p

63s

1(2.8.1)

inner shell 3s 3p

Magnesium Mg (Z = 12) 1s22s

22p

63s

2(2.8.2)

inner shell 3s 3p

Aluminium Al (Z = 13) 1s22s

22p

63s

23p

1(2.8.3)

inner shell 3s 3p

Silicon Si (Z = 14) 1s22s

22p

63s

23p

2(2.8.4)

inner shell 3s 3p

Phosphorus P (Z = 15) 1s22s

22p

63s

23p

3(2.8.5)

inner shell 3s 3p

Surphur S (Z = 16) 1s22s

22p

63s

23p

4(2.8.6)

inner shell 3s 3p

Chlorine Cl (Z = 17) 1s22s

22p

63s

23p

5(2.8.7)

inner shell 3s 3p

Argon Ar (Z = 18) 1s22s

22p

63s

23p

6(2.8.8)

inner shell 3s 3p

Note that : The 3s and 3p subshells of argon is completely filled. This electronic configuration is verystable. ( Argon is anoble gas )


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Unit 2-14

s-block elements of the fourth period :

Potassium K (Z = 19) 1s22s

22p

63s

23p

64s

1(2.8.8.1)

inner shell 3d 4s

Note that : The 4s orbital has energy lower than the 3d orbitals. According to Aufbau principle, the last

electron of potassium enters the 4s orbital, not the 3d orbital.

Calcium Ca (Z = 20) 1s22s

22p

63s

23p

64s

2(2.8.8.2)

inner shell 3d 4s

d-block elements of the fourth period : Transition elements ( Z = 21 30 )

Scandium Sc (Z = 21) 1s22s

22p

63s

23p

63d

14s

2(2.8.9.2)

inner shell 3d 4s

Note that : The last electron of scandium enters the 3d orbital. 4s orbital now has higher energy than 3d

orbitals.

Titanium Ti (Z = 22) 1s22s

22p

63s

23p

63d

24s

2(2.8.10.2)

inner shell 3d 4s

Vanadium V (Z = 23) 1s22s

22p

63s

23p

63d

34s

2(2.8.11.2)

inner shell 3d 4s

Chromium Cr (Z = 24) 1s22s

22p

63s

23p

63d

54s

1(2.8.13.1)

inner shell 3d 4s

Note that :

Chromium has the electronic configuration 1s22s

22p

63s

23p

63d

54s

1but not 1s

22s

22p

63s

23p

63d

44s

2.

The exactly half-filled 3d subshell and 4s subshell structures give extra stability to the chromium atom.

Manganese Mn (Z = 25) 1s22s

22p

63s

23p

63d

54s

2(2.8.13.2)

inner shell 3d 4s

Iron Fe (Z = 26) 1s22s

22p

63s

23p

63d

64s

2(2.8.14.2)

inner shell 3d 4s

Cobalt Co (Z = 27) 1s22s22p63s23p63d74s2 (2.8.15.2)inner shell 3d 4s


15/22


16/22

Unit 2-16

( 3 ) Electronic configurations in relation to the Periodic Table

In the Periodic Table, the elements are arranged in order ofincreasing atomic number. A new

row (period) is started when electrons start to enter a new principal energy level. Elements whose

atoms have a similar outer electronic configuration are placed in a vertical column (group). Within

any block of elements in the Periodic Table, the final electron to be added to an atom enters a subshell of

the type shown by the block letter (s, p, d or f).

Elements Descriptions-block Group I (alkali metals)

Group II (alkaline

earth metals)

1. All the elements are active metals.

2. Group similarities and trends within the groups are generally

clear.

p-block Group III to Group VI

Group VII (Halogens)

Group O (Noble gases)

1. Chemical behaviour in this block varies widely with the

reactivity of the metals, metalloids and non-metals, and the

comparative lack of reactivity of noble gases.2. Similarities within a group are shown by halogens and noble

gases.

3. Group IV elements carbon to lead show a dramatic change in

chemical properties within a group (a transition from

non-metal to metal).

d-block Transition elements 1. Transition elements are defined as elements which have

incomplete d subshell when combined in compounds.

2. They frequently have coloured compounds and complex ions

in various oxidation states.

3. They have higher melting points, and are denser and harder

than non-transition metals.f-block Lanthanoids and

Actinoids

1. Lanthanoids are widely spread throughout the earths crust.

They are also known as rare earth.

2. Actinoids are all radioactive metals.


17/22

Unit 2-17

Section 2.4 Electronic Structure and Ionisation Enthalpy

( 1 ) Ionisation enthalpy

Ionisation enthalpy is the energy absorbed when a mole of electron is removed from a mole of the

element in the gaseous state to give positively charged ions. Energy is always absorbed in this process

because work has to be done in overcoming the attractive force of the positively charged nucleus for the

negatively charged electrons. The energy is measured in kJ mol-1.

The first ionisation enthalpy of an element is defined as the energy required for the removal of one

mole of electrons from one mole of isolated atoms of the element in the gaseous state :

The second ionisation enthalpy is defined as the energy needed for the removal of a second mole of

electrons from one mole of isolated unipositive ions :

Higher ionisation enthalpy may be defined in a similar way. Thus, the nth

ionisation enthalpy is

the energy required for the process :

Example: kJ mol-1

at 298 K

I.E.1 I.E.2 I.E.3

sodium 500 4560 6940

magnesium 742 1450 7740

aluminium 583 1820 2740

Discussions:1. I.E.3> I.E.2> I.E.1

The positive charge on an ion increases as more electrons are removed. Also, as the charge on a

cation increases, its radius decreases. The nucleus then exerts a stronger electrostatic attractive force

on the subsequent electron to be removed.

2. For the alkali metals (e.g. sodium), thejump from I.E.1 to I.E.2is particularly large. The second

electron to be removed originates from a lower electron shell with stable noble gas electronic

configuration. It is closer to the nucleus and is very strongly held.

3. Similar effect is observed with elements of other groups where electrons are removed from a lower

electron shell. For example, a particularly large jump from I.E.2 to I.E.3 is observed for group II

elements (e.g. magnesium). The third electron to be removed originates from a lower electron shell

with stable noble gas electronic configuration.


18/22

Unit 2-18

( 2 ) Successive ionisation enthalpies for a particular element

The variation in the successive ionisation enthalpies for an element indicates the existence of

electron shells and subshells. Figure below shows the plots of successive ionisation enthalpies of a

potassium atom :

Interpretation of the plot :

1. The energy required to removed the 2nd to the 9th electrons, and the 10th to the 17th electrons,

increases steadily.

Interpretation : Each successive ionisation enthalpy is higher than the previous one because the positive

charge on an ion increases as more electrons are removed. The nucleus then exerts a stronger

electrostatic attractive force on the subsequent electron to be removed.

2. There are sharp and abrupt rises between the 1st

and 2nd

, 9th

and 10th

, and 17th

and 18th

ionisation

enthalpies.

Interpretation : Initially, the outer electrons which are loosely attracted to the nucleus are removed.

The electrons which are removed later are from lower energy levels and closer to the nucleus. The

last electrons are very strongly held to the nucleus.

Existence of electron shells :

1. In a potassium atom, the two electrons in the lowest energy level are closest to the nucleus. They are

most strongly attracted by the positive nucleus electrostatically. These two electrons are said to

occupy the first electron shell (n = 1).

2. The following 8 electrons occupy the second electron shell (n = 2). They have higher energy level

than the first 2 electrons do. They are also more distant from the nucleus.

3. Another 8 electrons, which are at a higher energy level, occupy the third electron shell (n = 3).

4. The last (outermost) electron occupies the fourth electron shell (n = 4). It has the highest energy level

and is most distant from the nucleus.


19/22

Unit 2-19

Section 2.5 The Periodic Table and atomic properties of elements

( 1 ) Trends of atomic radii of the elements in the Periodic Table

Figure below illustrates the variation of the atomic radius for the first 56 elements in the Periodic

Table :

The following trends can be observed :

1. On passing from left to right across a period, atomic radius of the elements decreases.

As the nuclear charge increases with atomic number, the additional electrons enter the same

electron shell. Within a period, the outermost electrons are poorly shielded from the attraction of

increasing nuclear charge by the outer electrons in the same shell. The effective nuclear charge ofsuccessive elements increases and the electrostatic attraction between outer electrons and positive

nucleus increases. Therefore, the outer electrons are drawn closer towards the nucleus, resulting in a

reduction in atomic radius.

Within a period, alkali metal atom has the largest size because the outer electron is in a new

electron shell, which is further away from the nuclear attraction.

The atom of noble gas has the smallest size because the electrons in the same shell do not shield

each other from the nuclear charge effectively while the effective nuclear charge is large.

2. The atomic radius increases down a groupElement has one more filled electron shell than the previous element in the same group. As the

electron occupy electron shells of greater quantum number in successive period, they become further

away from the nucleus.

Besides, the inner shell of electrons can shield the outer electrons effectively from the nuclear

charge. This decreases the attractive force between the positive nucleus and the electrons.

3. Across a period, the atomic radius of elements decrease abruptly at first but to the end of the

period,the extent of decrease becomes smaller.

At the beginning of each period, the added electrons experiences great effective nuclear charge

and the electron cloud contracts quickly.

Towards the end of a period, the effective nuclear charge increases comparatively slowly and the

corresponding decrease in atomic size is smaller.


20/22

Unit 2-20

( 2 ) Trends of the first ionisation enthalpies of the elements in the Periodic Table

The following figure illustrates the variation of the first ionisation enthalpy for the first 56 elements in the

Periodic Table :

The following trends can be observed :

1. On passing from left to right across each period, there is a general increase in the first ionisation

enthalpy.

As the nuclear charge increases with atomic number, the additional electrons enter the same

electron shell. In the same shell, electrons do not shield each other effectively from the positive

nuclear attraction. This results in an increase in effective nuclear charge. The outer electrons are

therefore attracted more tightly towards the nucleus. More energy is needed to remove an electron

from the atom.

2. Across each period, ionisation enthalpy peaks at Group O (Noble gases).

Noble gas electronic configurations are extremely stable because the ns and np subshells are

completely filled (ns2

np6). As the atomic radius of a noble gas is the smallest in a period, the

electrons are closet to the nucleus. Besides, the effective nuclear charge is highest and holds the

electrons tightly. More energy is needed to disturb the stable electronic structure and first ionisation

enthalpy is therefore extremely high.

3. There is a large drop in the first ionisation enthalpy in moving from one period to another.

At the beginning of each period, the additional s electron is added to a new electron shell further

away from the nucleus. Also, this s electron is effectively shielded from the attraction of nuclear

charge by the full inner shell of electrons. As a result, the outermost s electron is less firmly attractedand easily removed, leading to a particular low first ionisation enthalpy in Group I elements.

4. The first ionisation enthalpy decreases down a group.

The outer electrons are located one quantum shell further from the nucleus. As there are more

inner shell electrons, the outermost shell electrons are better shielded from the attraction of the nuclear

charge. Therefore, there is weaker attraction between the outer shell electrons and the nucleus. This

accounts for the decrease in first ionisation enthalpy down a group.


21/22

Unit 2-21

5. There is a drop in the first ionisation enthalpy between Group II and Group III ( Be and B;

Mg and Al).

Be has a fully-filled outermost 2s subshell. It is a stable electronic configuration and more

energy is required to remove an electron from it.

B has an additional 2p electron beyond the full 2s subshell. It is easier to remove an outermost

2p electron because this electron is further from the nucleus and is shielded from the attraction of the

nucleus by a full 2s subshell.

Mg has a fully-filled outermost 3s subshell. It is a stable electronic configuration and more

energy is required to remove an electron from it.

Al has an additional 3p electron beyond the full 3s subshell. It is easier to remove an outermost

3p electron because this electron is further from the nucleus and is shielded from the attraction of the

nucleus by a full 3s subshell.

6. There is a drop in the first ionisation enthalpy between Group V and Group VI ( N and O;

P and S).

N has an exactly half-filled outermost 2p subshell. This is a stable electronic structure because

the 2p orbitals are singly occupied and electronic charge is evenly distributed. Therefore electrostatic

repulsion is minimized. More energy is needed to remove a electron from this stable electronic

structure.

O has four electrons in the outermost 2p subshell. The fourth electron goes into an 2p orbital

which has already been occupied by an electron. The increased repulsion causes this electron to beless tightly held. Therefore, less energy is required to remove an outermost electron from O.

P has an exactly half-filled outermost 3p subshell. This is a stable electronic structure because

the 3p orbitals are singly occupied and electronic charge is evenly distributed. Therefore electrostatic

repulsion is minimized. More energy is needed to remove a electron from this stable electronic

structure.

S has four electrons in the outermost 3p subshell. The fourth electron goes into an 3p orbital

which has already been occupied by an electron. The increased repulsion causes this electron to be

less tightly held. Therefore, less energy is required to remove an outermost electron from S.


22/22

Unit 2-22

( 3 ) Trends of electron affinities of the elements in the Periodic Table

1. Electron affinity is the energy change when a mole ofgaseous atoms acquires a mole of electrons to

give a mole of negatively charged ions.

2. The general trend is for electron affinity to increase from left to right across the periodic table. This is

because, as the nuclear charge increases, the extra electrons enter the same electron shell and so are

attracted more strongly. Thus, atoms have progressively greater attraction for electrons as the group

number increases. Some electron affinities are given below, the values being in kJ mol-1

.

C

-120

N

0

O

-142

F

-333

S

-200

Cl

-348

Most electron affinities are exothermic because the orbitals of an electron theoretically extend to

infinity and so the nuclear charge is never completely balanced. However, when an electron is addedto an atom with a full or exactly half-filled outer subshell (e.g. nitrogen ), the electron affinity is low or

endothermic since these structures show enhanced stability. These values reinforce the deductions

made about stability from ionisation enthalpies.

3. The value usually decreases down a group of the Periodic Table as the atomic radii of the elements

increase. The reason is that the smaller the atom the closer is the outer electron shell to the nucleus

and the more strongly are electrons attracted into it.

F

-333

Cl-348

Br

-324

I

-295

The value for fluorine is anomalous because it includes a high repulsion electron cloud. This is

a consequence of a low atomic radius and a compact outer electron shell.

4. The second electron affinity of an element relates to the uptake of a second electron, and is defined as

the enthalpy required adding one mole of electrons to one mole of isolated uninegative ions.

The change is endothermic, because the repulsion component is greater than the attraction

component.

O S

First electron affinity -142 -200

Second electron affinity +791 +649

Electron affinities are generally obtained indirectly from thermochemical cycles because directmeasurement is difficult.

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Unit 2 Electronic