chemistry pre-u chemistry sem 1 chap 2

PRE-U STPM CHEMISTRY

SEMESTER 1

CHAPTER 2

ELECTRONIC STRUCTURE OF ATOMS

CHAPTER 2 : ELECTRONIC STRUCTURE OF ATOMS

2.1 Electronic Energy Levels of Atomic Hydrogen

2.2 Atomic orbitals

2.3 Electronic configuration

2.4Classification of elements into s, p, d and f blocks in the Periodic Table

Past Year Questions Analysis

Topic

2007 2008 2009 2010 2011 20122013

Sem 1

2014

Sem 1

P1 P2 P1 P2 P1 P2 P1 P2 P1 P2 P1 P2 AB

CA

B,

C

CHAPTER

2 :2 5a 2 1a 2 1 1 5a 1 1a 1 18 1 18

2.1 Electronic Energy Levels of Atomic Hydrogen

According to Maxwells theory, an electromagnetic wave has an

electric field component and a magnetic field component.

Electromagnetic radiation is the emission and transmission of energy

in the form of electromagnetic waves, which travel 3.00 x 108

meters per second (This constant is more well-known as speed of

light, c).,

Electromagnetic radiation is characterized by a frequency (f), and Electromagnetic radiation is characterized by a frequency (f), and

wavelength () where

=

=

cf;

)m(wavelength

)sm1000.3(lightofspeedFrequency

18

White light consists of continuous distribution of all possible

wavelengths spanning the entire visible region of the

electromagnetic spectrum. When a narrow beam of white light is

passed through a glass prism, different wavelengths travel

through the glass at different rates. As a result, the white light

dispersed into component colours, ranging from red at the long-

wavelength end of the spectrum (700 nm) to violet at the short-

wavelength end (400 nm)

When radiation from a particular source is passed through a

spectrometer, it will be separated into its components of

different frequencies, producing a spectrum.

i. A continuous spectrum is one where light is emitted over

a broad range of wavelengths (or frequencies); showing

emission of a wide range of energies. The spectrum is

smooth and continuous.

ii. A line spectrum is one where exact frequencies or

wavelengths appear as lines (indicating that only certainwavelengths appear as lines (indicating that only certain

amounts of energy are emitted and none in between those

energies). A line spectrum is normally produced by atoms

that have been excited and is also called atomic emission

spectrum

When a sample of hydrogen gas (H2) is subjected to an electrical

discharge, the hydrogen molecules dissociate forming hydrogen

atoms. Equation : H2 (g) H (g)

The hydrogen atoms formed, then absorb different amounts of

energy and the electron in each of the atoms will be raised to higher

energy level. Spectroscopists studying the spectrum of atomic

hydrogen had identified several series of spectral lines in different

regions of electromagnetic spectrum

a) These spectra line formed from spectroscopy are specific and

can be quantisised according to the radiation source by using

Rydbergs Equation

=

2

2

2

1

111

nnRH

= wavelength (in m)

n1 = ground state energy level

n2 = energy level where electron fall from

compared to ground state

RH = Rydberg constant = 1.097 x 107 m-1

b) Hydrogen spectrum produced under different sources produce

different series with different characteristics. Table 2.2 below

compared the hydrogen spectrum produced under ultraviolet

and visible ray

Ultraviolet rays Visible rays

Produced Lyman series Produced Balmer series

Usually use to calculate ionisation energy of

hydrogen gaseous atom

Can be used to determine wavelength produced by each

spectra given by dispersion of

light

Under emission spectrum, electrons from higher

energy level settled at n = 1

Under emission spectrum, electrons from higher energy

level settled at n = 2

Series of convergence lines produced have higher

frequency

Series of convergence lines produced have lower

frequency

2.1.1 The Bohr Model of the Hydrogen Atom

1. Neil Bohr suggested a model for the hydrogen atom that predicted the

existence of line spectra. He then outline 3 postulates about hydrogen

atom where :

The H atom has only certain allowable energy levels, which Bohr

called stationary states. Each of these states is associated with a fixed

circular orbit of the electron around the nucleus.

The atom does not radiate energy while in one of its stationary states

(atom does not change energy while the electron moves within an(atom does not change energy while the electron moves within an

orbit).

The atom changes to another stationary state (the electron moves to

another orbit) only by absorbing or emitting a photon whose energy

equals the difference in energy between the two states:

Estate A Estate B @ E =hf or E =hc /

(where the energy of state A is higher than that of state B, and h is the

constant proposed by Plancks theory where h = 6.63 x 10-34 J s.)

A spectral line results when a photon of specific energy (and thus

specific frequency) is emitted as the electron moves from a higher

energy state to a lower one.

Therefore, Bohrs model explains that an atomic spectrum is not

continuous because the atoms energy has only certain discretecontinuous because the atoms energy has only certain discrete

levels, or states

2. The quantum number, n (1, 2, 3, . . .) is associated with the shell

of an electron orbit, which is directly related to the electrons

energy: the lower the n value, the closer the orbit to the nucleus,

and the lower the energy level.

When the electron is in the first orbit (n = 1), the orbit closest to

the nucleus, the hydrogen atom is in its lowest energy level,

called the ground state.

If the hydrogen atom absorbs a photon whose energy equals

the difference between the first and second energy levels, thethe difference between the first and second energy levels, the

electron moves to the second orbit (n = 2), the next orbit further

from the nucleus. When the electron is in the second or any

higher orbit (energy level),the atom is said to be in an excited

state. The process where electron moved from ground state to

higher energy level is called as absorption

If the H atom in the first excited state emits a photon of that same

energy, it returns to the ground state. The process where electron

dropped from higher energy level (excited state) back to ground

state is called as emission

When a sample of gaseous Hydrogen atoms is excited, different

atoms absorb different quantities of energy. Each atom has one

electron, but so many atoms are present that all the energy levels

(orbits) are populated by electrons. When the electrons drop from

outer orbits to the n = 3 orbit (second excited state), the emitted

photons create the infrared series of lines. The visible series arises

when electrons drop to the n = 2 orbit (first excited state).

Ultraviolet

Lyman SeriesInfrared

Passchen

Series

Visible light

Balmer Series

When an electron drops from an outer orbit to an inner one, the

atom emits a photon of specific energy that gives rise to a spectral

line. In a given series, each electron drop, and thus each emission,

has the same inner orbit, that is, the same value of n1 in the

Rydberg equation, where the orbit radius is proportional to n2

value.

An energy diagram shows how the ultraviolet series arises. Within

each series the greater the difference in orbit radii, the greater the

difference in energy levels, and the higher the energy of the photon difference in energy levels, and the higher the energy of the photon

emitted. For example, in the ultraviolet series, in which n = 1, a

drop from n2 = 5 to n1 = 1 emits a photon with more energy

(shorter wavelength, higher frequency) than a drop from n2 = 2 to

n1 = 1. [The axis shows negative values because n = is defined as the atom with zero energy.]

Since Bohrs model is a one-electron model. It works beautifully for

the H atom and for other one-electron species, such as He+ (Z = 2),

Li2+ (Z = 3) and Be3+ (Z = 4),

One of the usefulness of Bohrs theory, applied when calculating

the energy levels of an atom, which he derived from the classical

principles of electrostatic attraction and circular motion, where the

equation is describe as

For hydrogen atom, since the atomic number, Z = 1. Therefore :

=

2

2181018.2

n

ZJE

If the ground level (under Lyman series) n = 1, the energy at ground

state is

=

2

18 11018.2n

JE

JEJE18

2

18 1018.2;1

11018.2 =

=

a. Note that even though the energy value is negative, however, as

mentioned above, under zero energy where E = 0 kJ when n = .In terms of magnitude, more energy will be released when

electron fall from n = to n = 1. If the ground state energy level ishigher, lesser energy will be released.

b. Derivation from equation above allowed us to find the energy

difference between two energy level, where

= @EEE

=

=

=

22

21

18

21

18

22

18

1n2n

n

1

n

1J1018.2E

n

1J1018.2

n

1J1018.2E

@EEE

c. Further derivation from equation (b) also allowed us to find the

wavelength produce in absorption / emission process. Using

Plancks equation, where

E = hf or E = hc /

=

22

21

18

substituteandrearrange

;hc

n

1

n

1J1018.2

=

=

22

21

7

22

21

834

18

n

1

n

11010.1

1

n

1

n

1

)1000.3)(1063.6(

J1018.21

substituteandrearrange

Ionisation energy of one mole of electron in hydrogen atom can

also be calculated using Bohrs equation.

H (g) H+ (g) + e- H = + x kJ mol-1 (Ionisation energy)

In order to remove an electron from hydrogen atom, electron must

at least reached the convergence limit of the energy level n = . Consider the ground state energy level of Lyman series, n = 1.

22

21

18 ;n

1

n

1J1018.2E

=

To remove 1 mol of electron from ground state, n1 = 1 to

convergence limit, n2= , a total of 1310 kJ is required.

1

3

2318

3A

18

22

18

molkJ1310H;J10

kJ1)1002.6(J1018.2H

;J10

kJ1NEH;emol1For

J1018.2E;1

1

1J1018.2E

==

=

=

=

Example 1 : Calculate the wavelength of the

first line of Lyman series


third line of Balmer series

Example 3 : Calculate the wavelength of the Example 4 : Calculate the wavelength of the

In Lyman series, n1 = 1

First line in series, so n2 = 1 + 1 = 2

=

2

2

2

1

111

nnRH

=

227

2

1

1

110097.1

1

= 122 nm

In Balmer series, n1 = 2


=

2

2

2

1

111

nnRH

=

227

5

1

2

110097.1

1

= 434 nm


forth line of Paschen series


last line of Balmer series

In Passchen series, n1 = 3


=

2

2

2

1

111

nnRH

=

227

7

1

3

110097.1

1

= 1005 nm


Last line in series, so n2 = 2 + =

=

2

2

2

1

111

nnRH

=

227 1

2

110097.1

1

= 365 nm

Example 5 : Calculate the frequency of the

second line of Paschen series

Example 6 : Calculate the frequency of the last

line of Lyman series


Second line in series, n2 = 3 + 2 = 5

=

2221

Hn

1

n

1R

1

= 711

10097.11


last line in series, so n2 = 1 + =

=

2

2

2

1

111

nnRH

= 711

10097.11

=

227

5

1

3

110097.1

1

1 / = 7.80 x 105 m-1

f = c x (1 / ) = (3.0 x 108)(7.80 x 105)

f = 2.34 x 1014 s-1

=

227 1

1

110097.1

1

1 / = 1.097 x 107 m-1

f =c x (1 / ) =(3.0 x 108)(1.097 x 107)

f = 3.29 x 1015 s-1

c) the third line of Lyman series d) the fifth line of Balmer series



=

2221

Hn

1

n

1R

1

= 711

10097.11


Fifth line in series, so n2 = 2 + 5 = 7

=

2

2

2

1

111

nnRH

= 711

10097.11

=

227

4

1

1

110097.1

1

1 / = 1.028 x 107 m-1

f = c x (1 / )

= (3.0 x 108)( 1.028 x 107)

f = 3.09 x 1015 s-1

=

227

7

1

2

110097.1

1

1 / = 2.519 x 107 m-1

f = c x (1 / )

= (3.0 x 108)( 2.519 x 107)

f = 7.56 x 1014 s-1

Example 3 : Using Bohrs Equation, calculate the energy required to cause the emission

of spectral line below

a) the fifth line of Lyman series

In Lyman series, n1= 1


E = 2.12 x 10-18 J / e-

b) the first line of Balmer series



E = 3.03 x 10-19 J / e-

=

2

2

2

1

18 111018.2nn

JE

=

22

18

6

1

1

11018.2 JE

=

2

2

2

1

18 111018.2nn

JE

=

22

18

3

1

2

11018.2 JE

E = 2.12 x 10-18 J / e- E = 3.03 x 10-19 J / e-

c) the third line of Passchen series



E = 1.82 x 10-19 J / e-

d) the last line of Passchen series


Last line in series, so n2 = 3 + =

E = 2.42 x 10-19 J / e-

=

2

2

2

1

18 111018.2nn

JE

=

22

18

6

1

3

11018.2 JE

=

2

2

2

1

18 111018.2nn

JE

=

22

18 1

3

11018.2 JE

2.2 Atomic Orbital

The position of electrons cannot be specified as electron behaves like

wave as it extended it space. Werner Karl Heisenberg, then formulated

what is now known as the Heisenberg uncertainty principle: it is

impossible to know simultaneously both the momentum and the position

of a particle with certainty.

However, Bohrs theory had made a significant contribution to our

understanding of atoms, and his suggestion that the energy of an

electron in an atom is quantized. This concept is the perfected by an

Austrian physicist, Erwin Schrdinger, through his well-knownAustrian physicist, Erwin Schrdinger, through his well-known

equation Schrdingers equation, where the energy of atom can be

calculate.

Even though Schrdinger equation specifies the possible energy states

the electron can occupy in a hydrogen atom, however, it cannot pin-

point the location of electron in an atom. Therefore, to counter this

problem, we replaced with the term orbital, a region with high

probability to find an electron.

An atomic orbital is specified by three quantum numbers. One is related

to the orbitals size, another to its shape, and the third to its orientation

in space. The quantum numbers have a hierarchical relationship: the

size-related number limits the shape-related number, which limits the

orientation-related number

The principal quantum number (n) (better known as shell) is a positive integer (1, 2, 3, and so forth). It indicates the relative size of the orbital

and therefore the relative distance from the nucleus of an atom.

The angular momentum quantum number(l) is an integer from 0 to n

1. It is related to the shape of the orbital and is sometimes called the 1. It is related to the shape of the orbital and is sometimes called the

orbital shape (or azimuthal) quantum number. Note that the principal

quantum number sets a limit on the values for the angular momentum

quantum number; that is, n limits l.

For an orbital (shell) with n = 1, l can have a value of only 0.

For orbitals (shell) with n = 2, l can have a value of 0 or 1

For orbitals (shell) with n = 3, l can be 0, 1, or 2; and so forth.

Note that the number of possible l values equals the value of n

The magnetic quantum number (ml) is an integer from [l] through

0 to [+l]. It prescribes the orientation of the orbital in the space

around the nucleus (or simple, number of orbitals presence in l).

The possible values of an orbitals magnetic quantum number are

set by its angular momentum quantum number.

For (l = 0), magnetic quantum number, (ml) = 0 [therefore 1 orbital]

For (l = 1), magnetic quantum number, (ml) = 1, 0, +1 [therefore 3

orbitals]

For (l = 2), magnetic quantum number, (m ) = 2, 1, 0, +1, +2 For (l = 2), magnetic quantum number, (ml) = 2, 1, 0, +1, +2

[therefore 5 orbitals]

The electron spin quantum number (ms) ~ represents the

assumption of electrons act like tiny magnets. According to

electromagnetic theory, a spinning charge generates a magnetic

field, and it is this motion that causes an electron to behave like a

magnet. Therefore, in each ml, two oppositely spin quantum is filed

accordingly and has a value of + and , and are usually denote as

(for +) and (for )

The energy states and orbitals of the atom are described with

specific terms and associated with one or more quantum numbers

Level. The atoms energy levels, or shells, are given by the n

value: the smaller the n value, the lower the energy level and

the greater the probability of the electron being closer to the

nucleus. [n = 1 is the closest to nucleus, followed by n = 2, 3

and so forth]

Sublevel. The atoms levels contain sublevels, or subshells,

which designate the orbital shape. Each sublevel has a letter which designate the orbital shape. Each sublevel has a letter

designation:

l 0 1 2 3 4 5

Name of sublevel

(orbital)s p d f g h

Orbital. Each allowed combination of n, l, and ml values specifies

one of the atoms orbitals. Thus, the three quantum numbers that

describe an orbital express its size (energy), shape, and spatial

orientation. You can easily give the quantum numbers of the

orbitals in any sublevel if you know the sublevel letter designation

and the quantum number hierarchy

Energy

level, n

Sub-

level, lorbital, m

l

No of

orbital

Atomic Orbital

Designation

10 0 1 1s

20

1

0

1, 0 +1

1

3

2s

2px , 2py , 2pz,

3

0

1

2

0

1, 0 +1

2, 1, 0 , +1 , +2

1

3

5

3s

3px , 3py , 3pz,

3dxy , 3dyz , 3dxz , 3dx2-y2 , 3dz2

2. Shape of each orbitals

The s Orbital An orbital with l = 0 has a spherical shape with the

nucleus at its center and is called an s orbital

The p orbitals An orbital with l = 1, called a p orbital, has two The p orbitals An orbital with l = 1, called a p orbital, has two

regions (lobes)of high probability, one on either side of the nucleus.

The nucleus lies at the nodal plane of this dumbbell-shaped orbital

as described in diagram below

The d Orbital An orbital with l = 2 is called a d orbital. There are

five possible ml values for the l = 2 value: 2, 1, 0 , +1 , +2. Thus, a

d orbital can have any one of five different orientations, as describe

in diagram below

Orbitals with Higher l values Orbitals with l = 3 are f orbitals and

must have a principal quantum number of at least n = 4. There are

seven f orbitals (2 l + 1 = 7), each with a complex, multi-lobed shape

Special case of hydrogen atom The

energy state of the H atom depends only

on the principal quantum number, n = 1.

When an electron occupies an orbital with

a higher n value, it occurs from the

nucleus, so the atom is higher in energy.

But the H atom is a special case because it

has only one electron. The energy states of

all other atoms depend on both the n and l

values of the occupied orbitals because ofvalues of the occupied orbitals because of

additional nucleus electron attractions

and electron electron repulsions. In

other words, for the H atom only, all four

orbitals in n = 2 (one 2s and three 2p) have

the same energy, while all nine orbitals in

n = 3 (one 3s, three 3p, and five 3d) have

the same energy and so forth.

2.3 Electronic configuration

Electron configuration of the atom shows how the electrons are

distributed among the various atomic orbitals, in order to

understand electronic behaviour of that atom. Using the principle

of n, l, ml and ms learned earlier, it allows us to understand how the

arrangement of electrons occurs in many-electrons atom.

The arrangement of electrons in its orbitals are guided under 3

basic rule and principles, which are Aufbau's Principle, Pauli

Exclusion's Principle and Hund's Rule.Exclusion's Principle and Hund's Rule.

Aufbau's principle stated that electrons are filled up in orbitals

from the lowest energy orbital available. This will results in ground-

state electron configurations to build up eventually

Pauli Exclusion's Principle ~ an atomic orbital can hold a

maximum of two electrons with opposing quantum spins. From

the quantum spin number, we understand that electrons

behaviour resemble to that of a magnet when spinning charge

generates a magnetic field. In general, we represent a positive

spin quantum, ms = +1/2 as (sometimes ) ; while a negativespin quantum, ms = -1/2 as (sometimes ).

Correct IncorrectCorrect Incorrect

2 electrons occupied an

orbital with opposite spin

2 electrons occupied

an orbital with same

spin


an orbital with same

spin


an orbital with

different spin

Hund's Rule ~ when orbitals of equal energy are available, the

electron configuration of lowest energy has the maximum number

of unpaired electrons with parallel spins. In order to fulfil Hund's

rule, sub-shell must have at least 2 or more orbitals. Therefore, p-

orbitals, d-orbitals and f-orbitals are filled according to Hund's rule.

For example, in filling 2 and 3 electrons in p-orbitals and filling 5

and 7 electrons in d-orbitals

Filling in p - orbitals Filling in d - orbitals

Filling in 2 electrons in p-orbitals Filling in 5 electrons in d-orbitals (more stable)

Filling in 3 electrons in p-orbitals

(more stable)

Filling in 7 electrons in d-orbitals (After positive spin is first

filled, negative spin is then filled to each orbital)

ElementNo

of e-Orbital diagram

Electronic

configuration

Hydrogen

H____

1s

Helium

He____

1s

Lithium

Li ____ ____

1s 2s

1 1s1

2 1s2

3 1s22s1

Beryllium

Be ____ ____

1s 2s

Boron

B____ ____ ____ ____ ____

1s 2s 2 p

Carbon

C____ ____ ____ ____ ____

1s 2s 2 p

Nitrogen

N____ ____ ____ ____ ____

1s 2s 2 p

4 1s22s2

5 1s22s22p1

6 1s22s22p2

7 1s22s22p3

ElementNo

of e-Orbital diagram

Electronic

configuration

Oxygen

O____ ____ ____ ____ ____

1s 2s 2 p

Fluorine

F____ ____ ____ ____ ____

1s 2s 2 p

Neon

Ne____ ____ ____ ____ ____

1s 2s 2 p

Sodium____ ____ ____ ____ ____ ____

8 1s22s22p4

9 1s22s22p5

10 1s22s22p6

11 1s22s22p63s1Sodium

Na____ ____ ____ ____ ____ ____

1s 2s 2 p 3s

Magnesium

Mg____ ____ ____ ____ ____ ____

1s 2s 2 p 3s

Aluminium

Al____ ____ ____ ____ ____ ____ ___ ___ ___

1s 2s 2 p 3s 3 p

Silicon

Si____ ____ ____ ____ ____ ____ ____ ____ ____

1s 2s 2 p 3s 3 p

11 1s22s22p63s1

12 1s22s22p63s2

13 1s22s22p63s23p1

14 1s22s22p63s23p2

ElementNo of

e-Orbital diagram Electronic configuration

Phospho-

rous, P____ ____ ____ ____ ____ ____ ____ ____ ____

1s 2s 2 p 3s 3 p

Sulphur

S____ ____ ____ ____ ____ ____ ____ ____ ____

1s 2s 2 p 3s 3 p

Chlorine

Cl____ ____ ____ ____ ____ ____ ____ ____ ____

1s 2s 2 p 3s 3 p

15 1s22s22p63s23p3

16 1s22s22p63s23p4

17 1s22s22p63s23p5Cl1s 2s 2 p 3s 3 p

Argon

Ar____ ____ ____ ____ ____ ____ ____ ____ ____

1s 2s 2 p 3s 3 p

Potassium

K____ ____ ____ ____ ____ ____ ____ ____ ____ ___

1s 2s 2 p 3s 3 p 4s

Calcium

Ca____ ____ ____ ____ ____ ____ ____ ____ ____ ___

1s 2s 2 p 3s 3 p 4s

18 1s22s22p63s23p6

19 1s22s22p63s23p64s1

20 1s22s22p63s23p64s2

1s22s22p63s23p63d14s2

1s22s22p63s23p63d24s2

1s22s22p63s23p63d34s2

1s22s22p63s23p63d54s1

1s22s22p63s23p63d54s2

1s22s22p63s23p63d64s2

1s22s22p63s23p63d74s2

1s22s22p63s23p63d84s2

1s22s22p63s23p63d104s1

1s22s22p63s23p63d104s2

a. Note that from Scandium to Vanadium, each electron is filled

according to Hund's rule, with a single positive spin electron is

filled in each 3d-subshells.

b. When expressing the electronic configuration for Chromium, 24Cr,

the valence electron of Cr is filled as 3d5 4s1 instead of 3d4 4s2.

c. This is due to, according to Hund's rule, half-filled 3d orbitals have

extra stability, compared to a partial-filled 3d orbital.

Another anomaly of filling the electronic configuration occur on the

element copper, Cu. Supposedly, After nickel, 28Ni is filled as

1s22s22p63s23p63d84s2, Cu should be filled : 1s22s22p63s23p63d94s2.

However, due to full-filled 3d orbitals have extra stability

compared to a partial-filled 3d orbitals, henceforth valence

electrons of Cu is filled as 3d104s1.

2.3.1 Electronic Configuration of Ions

1. Ions are formed when an atom or molecule donate / received

electron(s). Ions can be positively charged or negatively charged. A

positively charged ion is also known as cation, while a negatively

charged ion is also known as anion. Table below compared the

properties of the formation for both cation and anion

Ions Cation (Positively charged ion) Anion (Negatively charged ion)

Occur when Electron(s) are donated Electron(s) are received

Formation of

+1 and -1

Na Na+ + e-

1s22s22p63s1 1s22s22p6F + e- F-

1s22s22p5 1s22s22p6

Formation of

+2 and -2

Mg Mg2+ + 2e-

1s22s22p63s2 1s22s22p6O + 2e- O2-

1s22s22p4 1s22s22p6

2. From the example above, it is shown that, when electron(s) are

donated, electron(s) are first removed from higher energy level,

and conversely when electron(s) are received, electron(s) are filled

from the lower possible energy level. Most of the main group

elements donate and received electron(s) to achieve a stable

valence electronic configuration of ns2 np6 (also known as octet

configuration)

Al3+ : P3-1s22s22p6 1s22s22p63s23p6

[Since electronic configuration of

aluminium is 1s22s22p63s23p1, when 3

electrons are removed, it shall be removed

from 3p, then 3s]


aluminium is 1s22s22p63s23p3, when 3

electrons are added, it shall be added to 3p

orbital as its not yet complete]

V3+


vanadium is 1s22s22p63s23p63d34s2 when 3

electrons are removed, it shall be removed

from 4s, then 3d]

N3-

[Since electronic configuration of nitrogen

is 1s22s22p3, when 3 electrons are added, it

shall be added to 3p orbital as its not yet

complete]

1s22s22p63s23p63d2 1s22s22p6

Fe2+

[Since electronic configuration of iron is

.when 2

electrons are removed, it shall be

removed from .]

S4-

[Since electronic configuration of sulphur

is .,,,,,,,,,,,,, when 4 electrons

are added, it shall be added to

. and .. orbitals]

Ga4+

[Since electronic configuration of gallium

is ... when

Br-

[Since electronic configuration of Br is

,

1s22s22p63s23p63d64s2

4s

1s22s22p63s23p63d6

1s22s22p63s23p4

3p 4s

1s22s22p63s23p64s2

1s22s22p63s23p63d104s24p1

1s22s22p63s23p63d9

1s22s22p63s23p63d104s24p5

1s22s22p63s23p63d104s24p6

is ... when

4 electrons are removed, it shall be

removed from ..]

,

when 1 electrons are added, it shall be

added to . orbitals]

Mn4+ Cl-

Ca+ O3-

Co3+ F-

1s22s22p63s23p63d104s24p1

4p , 4s , then 3d

1s22s22p63s23p63d104s24p5

4p

1s22s22p63s23p63d3 1s22s22p63s23p6

1s22s22p63s23p64s1 1s22s22p63s1

1s22s22p63s23p63d6 1s22s22p6

2.4 Classification of elements into s, p, d and f blocks in the Periodic

Table

chemistry pre-u chemistry sem 1 chap 2

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