-
1 Major Paper - 1 General Chemistry – 1 First - Year Major Paper
- I General Chemistry – I Unit 1: Atomic Structure, Periodic
Properties, Theories Of Volumetric And Semimicro Qualitative
Analyses. 1.1 Atomic Orbitals, quantum numbers - Principal,
azimuthal, magnetic and spin quantum numbers and their
significance-Principles governing the occupancy of electrons in
various quantum levels-Pauli’s exclusion principle, Hund’s rule,
Aufbau Principle, (n+l) rule, stability of half-filled and fully
filled orbitals. 1.2 Classification as s, p, d & f block
elements, variation of atomic volume, atomic and ionic radii,
ionisation potential, electron affinity and electronegativity along
periods and groups – Factors influencing the periodic properties.
1.3 Inorganic Qualitative Analysis: Solubility Product – Principle
of Elimination of interfering anions, Common ion Effect –
Complexation reactions including spot tests in qualitative analysis
– Reactions involved in separation and identification of cations
and anions in the analysis – Semi Micro Techniques . 1.4
Titrimetry: Definitions of Molarity, normality, molality and mole
fraction – Primary and Secondary standards – Types of titrimetric
reactions – acid-base, redox, precipitation and complexometric
titrations – Indicators – Effect of change in pH – theory of
neutralization, redox, adsorption and metal ion indicators. Unit 2:
Chemical bonding and chemistry of s- block elements 2.1 lonic bond
– Lattice Energy – Born – Haber Cycle – Pauling and Mulliken’s
scales of electronegativity – Polarizing power and Polarisability –
partial ionic character from electro- negativity – Transitions from
ionic to covalent character and vice versa – Fajan’s rules. 2.2
VESPR Theory – Shapes of simple inorganic molecules (BeCl2, SiCl4,
PCl5, SF6, IF7, NH3, XeF4, XeF6, XeO3, XeOF4, BF3, and H2O ) - VB
Theory – Principles of hybridization – MO Theory – Bonding and
antibonding orbitals – Application of MO Theory to H2, He2, N2, O2,
HF and CO – Comparison of VB and MO Theories. 2.3 S-block elements
General characteristics of Group IA elements – diagonal
relationship between Li and Mg – general methods of extraction,
physical, chemical properties and uses of Lithium, Sodium and
Potassium.
-
2 General characteristics of Elements of Group IIA – diagonal
relationship between Be and AI – Extraction of Beryllium, Magnesium
and Calcium – Physical, chemical properties and Uses. Unit 3:
Theories of Covalent Bonding and Structure, Chemistry of Alkanes
and Cycloalkanes 3.1 Covalent bonding – hybridization – Structure
of organic molecules based on sp³, sp² and sp hybridization
–properties of covalent molecules: bond length, bond angle, bond
energy, bond polarity, inductive, mesomeric, electromeric,
resonance and hyperconjugative effects. 3.2 IUPAC nomenclature of
organic compounds (up to 10 carbon systems) – Hydrocarbons – Mono
functional compounds – Bifunctional compounds – Isomerism – Types
of isomerism (structural and stereoisomerisms) with appropriate
examples. 3.3 Petroleum- Coal tar distillation- source of alkanes –
Methods of preparing alkanes and cycloalkanes - Chemical properties
– Halogenation of alkanes - Mechanism of free radical substitution
– Conformational study of ethane and n-butane - Relative stability
of cycloalkanes (from cyclopropane up to cyclooctane) – Bayer’s
Strain theory – Limitations – Conformational study of cyclohexane,
mono and disubstituted cyclohexanes. Unit 4: Chemistry of Alkenes,
Alkynes, Dienes, Benzene and Benzenoid Compounds. 4.1 Chemistry of
alkenes, alkynes and dienes: Nomenclature of alkenes – Geometrical
Isomerism –– General methods of preparation of alkenes – Chemical
properties – Uses – Elimination mechanisms (E1,E2) – Electrophilic,
Free radical additions – Ziegler – Natta Catalytic polymerization
of ethylene – polymers of alkene derivatives. Nomenclature of
alkynes- General methods of preparation of alkynes – Physical and
Chemical properties – Uses. 4.2 General methods of preparation of
Dienes and types- Physical and Chemical properties – Uses –
Mechanisms of electrophilic and Free radical addition reactions –
Rubber as a natural polymer. 4.3 General methods of preparation of
benzene – Chemical properties – Uses – Electrophilic substitution
mechanism – Orientation and reactivity in substituted benzenes.
Polynuclear Aromatic compounds – Nomenclature –Laboratory
preparation of Naphthalene, Anthracene and Phenanthrene -Structure,
aromatic character, physical, chemical properties and uses of
Naphthalene, Anthracene and Phenanthrene – Mechanism of Aromatic
electrophilic substitution – Theory of orientation and reactivity
Unit 5: Atomic Structure, Molecular Velocities and Real Gases 5.1
Dualism of light – Wave nature of radiation, classical theory of
electromagnetic radiation and classical expression for energy in
terms of amplitude. Particle nature of radiation – Black body
radiation and Planck’s quantum theory, photoelectric effect and
Compton effect of matter – de Broglie hypothesis and Davisson and
Germer experiment. Heisenberg’s uncertainty principle. Schrodinger
wave equation – arguments in favour of Schrodinger wave
-
3 equation. Physical significance of ψ and ψ2 functions. Wave
picture of electron – Concept of atomic orbitals. Shapes of s, p
and d orbitals. Nodal planes and nodal points in atomic orbitals- g
and u character of atomic orbitals. 5.2 Maxwell’s distribution of
Molecular velocities (Derivation not required). Types of Molecular
velocities – Mean, Most probable and root mean square velocities.
Graphical representation and its significance – Collision diameter,
Mean free path and collision number. Real gases – van der Waals
equation of states – derivation.Law of corresponding states.
Recommended Books: 1. Soni. P.L. Textbook of Inorganic Chemistry,
S. Chand & sons, New Delhi, (2010) 2. Soni. P.L. and Chawla
H.M., Textbook of Organic Chemistry, S. Chand & sons, New
Delhi, (2010) 3. Soni. P.L. and Dharmarha O.P. Textbook of Physical
Chemistry, S. Chand & sons, New Delhi, (2010) 4. Puri
B.R..,Sharma L.R., Kalia K.K., Principles of Inorganic Chemistry,
(23rd edition), New Delhi, Shoban Lal Nagin Chand & Co.,
(1993). 5. Puri B.R..,Sharma L.R., Pathania M.S., Principles of
Physical Chemistry, (23rd edition), New Delhi, Shoban Lal Nagin
Chand & Co., (1993). Books for reference: 1. Lee J.D., Concise
Inorganic Chemistry, UK, Black well science (2006). 2. Glasstone
S., Lewis D., Elements of Physical Chemistry, London, Mac Millan
& Co. Ltd. 3. Morrison R.T. and Boyd R.N., Organic Chemistry
(6th edition), New York, Allyn & Bacon Ltd., (1976). 4. Bahl
B.S. and Arun Bahl, Advanced Organic Chemistry, (12th edition), New
Delhi, Sultan Chand & Co., (1997). 5. Frank J. Welcher and
Richard B. Hahn, Semi micro Qualitative Analysis, New Delhi,
Affiliated East-west Press Pvt.Ltd. (1969).
-
Unit 1: Atomic Structure, Periodic Properties, Theories
Qualitative Analyses. 1.1. Atomic orbitals Atomic orbitals are
regions of space around the nucleus of an atom where an electron is
likely to be found. Atomic orbitals allow atoms to commonly filled
orbitals are s,spherical. P orbitals have a single angular node
across the nucleus and are shaped like dumbbells. D and f have two
and three angular nodes,be found in any orbital space as defined by
the Pauli orbitals is completely filled, a new set of orbitals
(have one angular node, like 1p, but also one radial node.1.1.1.
Orbitals The number of orbitals in a subshell is equivalent to the
number of values the magnetic quantum number ml takes on. A helpful
equation to determine the number of orbitals in a subshell is 2l
+1. This equation will not give you the value of ml, but the number
of possible values that ml can take on in a particular orbital. For
example, if l=1 and ml have values -1, 0, or +1, the value of 2l+1
will be three and there will be three different orbitals. The names
of the orbitals are named after the subshells they are found in: In
the figure below, we see examples of two orbitals: the p orbital
(blue)orbital (red). The red s orbital is a 1s orbital. To picture
a 2s orbital, imagine a layer similar to a cross section of a
jawbreaker around the circle. The layers are depicting the atoms
angular nodes. To picture a 3s orbital, imagine another lThe p
orbital is similar to the shape of a dumbbell, with its orientation
within a subshell depending on ml. The shape and orientation of an
orbital depends on l and m4 mic Structure, Periodic Properties,
Theories of Volumetric and Atomic orbitals are regions of space
around the nucleus of an atom where an electron is likely to be
found. Atomic orbitals allow atoms to make covalent bonds. The most
s, p, d, and f. S orbitals have no angular nodes and are orbitals
have a single angular node across the nucleus and are shaped like
have two and three angular nodes, respectively. Only two electrons
will be found in any orbital space as defined by the Pauli
Exclusion Principle. Once a shell of orbitals is completely filled,
a new set of orbitals (2s, 2p, and so on) is available.but also one
radial node. in a subshell is equivalent to the number of values
the m number ml takes on. A helpful equation to determine the
number of orbitals in a subshell is 2l +1. This equation will not
give you the value of ml, but the number of possible values that ml
can take on in a particular orbital. For example, if l=1 and ml 1,
0, or +1, the value of 2l+1 will be three and there will be three
different orbitals. The names of the orbitals are named after the
subshells they are found in:In the figure below, we see examples of
two orbitals: the p orbital (blue)orbital (red). The red s orbital
is a 1s orbital. To picture a 2s orbital, imagine a layer similar
to a cross section of a jawbreaker around the circle. The layers
are depicting the atoms angular nodes. To picture a 3s orbital,
imagine another layer around the circle, and so on and so on. The p
orbital is similar to the shape of a dumbbell, with its orientation
within a subshell . The shape and orientation of an orbital depends
on l and ml. nd Semi micro Atomic orbitals are regions of space
around the nucleus of an atom where an electron make covalent
bonds. The most orbitals have no angular nodes and are orbitals
have a single angular node across the nucleus and are shaped like
respectively. Only two electrons will . Once a shell of and so on)
is available. 2p would in a subshell is equivalent to the number of
values the m number ml takes on. A helpful equation to determine
the number of orbitals in a subshell is 2l +1. This equation will
not give you the value of ml, but the number of possible values
that ml can take on in a particular orbital. For example, if l=1
and ml can 1, 0, or +1, the value of 2l+1 will be three and there
will be three different orbitals. The names of the orbitals are
named after the subshells they are found in: In the figure below,
we see examples of two orbitals: the p orbital (blue) and the s
orbital (red). The red s orbital is a 1s orbital. To picture a 2s
orbital, imagine a layer similar to a cross section of a jawbreaker
around the circle. The layers are depicting the atoms angular ayer
around the circle, and so on and so on. The p orbital is similar to
the shape of a dumbbell, with its orientation within a subshell
-
5 1.1.2 Principal quantum number In quantum mechanics, the
principal quantum number (symbolized n) is one of four quantum
numbers which are assigned to all electrons in an atom to describe
that electron's state. As a discrete variable, the principal
quantum number is always an integer. As n increases, the number of
electronic shells increases and the electron spends more time
farther from the nucleus. As n increases, the electron is also at a
higher energy and is, therefore, less tightly bound to the nucleus.
The total energy of an electron, as described below, is a negative
inverse quadratic function of the principal quantum number n. The
principal quantum number was first created for use in the
semiclassical Bohr model of the atom, distinguishing between
different energy levels. With the development of modern quantum
mechanics, the simple Bohr model was replaced with a more complex
theory of atomic orbitals. However, the modern theory still
requires the principal quantum number. Apart from the principal
quantum number, the other quantum numbers for bound electrons are
the azimuthal quantum number ℓ, the magnetic quantum number ml, and
the spin quantum number s. There are a set of quantum numbers
associated with the energy states of the atom. The four quantum
numbers n, ℓ, m, and s specify the complete and unique quantum
state of a single electron in an atom, called its wave function or
orbital. Two electrons belonging to the same atom cannot have the
same values for all four quantum numbers, due to the Pauli
Exclusion Principle. The wave function of the Schrödinger wave
equation reduces to the three equations that when solved lead to
the first three quantum numbers. Therefore, the equations for the
first three quantum numbers are all interrelated. The principal
quantum number, n, designates the principal electron shell. Because
n describes the most probable distance of the electrons from the
nucleus, the larger the number n is, the farther the electron is
from the nucleus, the larger the size of the orbital, and the
larger the atom is. n can be any positive integer starting at 1, as
n=1n=1 designates the first principal shell (the innermost shell).
The first principal shell is also called the ground state, or
lowest energy state. This explains why n cannot be 0 or any
negative integer, because there exists no atoms with zero or a
negative amount of energy levels/principal shells. When an electron
is in an excited state or it gains energy, it may jump to the
second
-
6 principle shell, where n=2n=2. This is called absorption
because the electron is "absorbing" photons, or energy. Known as
emission, electrons can also "emit" energy as they jump to lower
principle shells, where n decreases by whole numbers. As the energy
of the electron increases, so does the principal quantum number,
e.g., n = 3 indicates the third principal shell, n = 4 indicates
the fourth principal shell, and so on. n=1,2,3,4… 1.1.3 Azimuthal
Quantum Number The azimuthal quantum number is a quantum number for
an atomic orbital that determines its orbital angular momentum and
describes the shape of the orbital. The azimuthal quantum number is
the second of a set of quantum numbers which describe the unique
quantum state of an electron (the others being the principal
quantum number, following spectroscopic notation, the magnetic
quantum number, and the spin quantum number). It is also known as
the orbital angular momentum quantum number, orbital quantum number
or second quantum number, and is symbolized as ℓ. Connected with
the energy states of the atom's electrons are four quantum numbers:
n, ℓ, mℓ, and ms. These specify the complete, unique quantum state
of a single electron in an atom, and make up its wave function or
orbital. The wave function of the Schrödinger equation reduces to
three equations that when solved, lead to the first three quantum
numbers. Therefore, the equations for the first three quantum
numbers are all interrelated. The azimuthal quantum number arose in
the solution of the polar part of the wave equation as shown below.
To aid understanding of this concept of the azimuth, it may also
prove helpful to review spherical coordinate systems, and/or other
alternative mathematical coordinate systems besides the Cartesian
coordinate system. Generally, the spherical coordinate system works
best with spherical models, the cylindrical system with cylinders,
the cartesian with general volumes, etc.
-
Atomic orbitals have distinctive shapes denoted by letters. In
the illustration, the letters s, p, and d describe the shape of
theTheir wavefunctions take the form ofby Legendre polynomials. The
various orbitals relating to different values ofsometimes called
sub-shells, and (mainly for historical reasons) are referred to by
letters, as follows: Quantum Subshells for the Azimuthal Quantum
NumberAzimuthal number (ℓ) Historical Letter MaximumElectrons0 s 2
1 p 6 2 d 10 7 Illustration of quantum mechanical orbital angular
momentum.Atomic orbitals have distinctive shapes denoted by
letters. In the illustration, the describe the shape of the atomic
orbital. take the form of spherical harmonics, and so are described
. The various orbitals relating to different values of, and (mainly
for historical reasons) are referred to by letters, Quantum
Subshells for the Azimuthal Quantum Number Maximum Electrons
Historical Name Shape sharp spherical principal three
dumbbell-shaped polar-aligned orbitals; one lobe on each pole of
the x, y, and z axes); two electrons each lobe. diffuse nine
dumbbells and one doughnut (or “unique shape #1” see this picture
of spherical harmonics, third row center) Illustration of quantum
mechanical orbital angular momentum. Atomic orbitals have
distinctive shapes denoted by letters. In the illustration, the ,
and so are described . The various orbitals relating to different
values of ℓ are , and (mainly for historical reasons) are referred
to by letters, aligned orbitals; pole of the x, y, and z (+ and −
doughnut (or “unique this picture of spherical
-
8 3 f 14 fundamental “unique shape #2” (see this picture of
spherical harmonics, bottom row center) 4 g 18 5 h 22 6 i 26 The
letters after the f sub-shell just follow letter f in alphabetical
order except the letter j and those already used. Each of the
different angular momentum states can take 2(2ℓ + 1) electrons.
This is because the third quantum number mℓ (which can be thought
of loosely as the quantizedprojection of the angular momentum
vector on the z-axis) runs from −ℓ to ℓ in integer units, and so
there are 2ℓ + 1 possible state. Each distinct n, ℓ, mℓ orbital can
be occupied by two electrons with opposing spins (given by the
quantum number ms = ±½), giving 2(2ℓ + 1) electrons overall.
Orbitals with higher ℓ than given in the table are perfectly
permissible, but these values cover all atoms so far discovered.
For a given value of the principal quantum number n, the possible
values of ℓ range from 0 to n − 1; therefore, the n = 1 shell only
possesses an s subshell and can only take 2 electrons, the n = 2
shell possesses an s and a p subshell and can take 8 electrons
overall, the n = 3 shell possesses s, p, and d subshells and has a
maximum of 18 electrons, and so on. Generally speaking, the maximum
number of electrons in the nth energy level is 2n2. The angular
momentum quantum number, ℓ, governs the number of planar nodes
going through the nucleus. A planar node can be described in an
electromagnetic wave as the midpoint between crest and trough,
which has zero magnitude. In an s orbital, no nodes go through the
nucleus, therefore the corresponding azimuthal quantum number ℓ
takes the value of 0. In a p orbital, one node traverses the
nucleus and therefore ℓ has the value of 1. Depending on the value
of n, there is an angular momentum quantum number ℓ and the
following series. The wavelengths listed are for a hydrogen atom: �
Lyman series (ultraviolet) � Balmer series (visible)
-
9 � Ritz-Paschen series (near infrared) � Brackett series
(short-wavelength infrared) � Pfund series (mid-wavelength
infrared). 1.1.3.1 The Orbital Angular Momentum Quantum Number (ℓ)
The orbital angular momentum quantum number ℓ determines the shape
of an orbital, and therefore the angular distribution. The number
of angular nodes is equal to the value of the angular momentum
quantum number ℓ. (For more information about angular nodes, see
Electronic Orbitals.) Each value of l indicates a specific s, p, d,
f subshell (each unique in shape.) The value of l is dependent on
the principal quantum number n. Unlike n, the value of l can be
zero. It can also be a positive integer, but it cannot be larger
than one less than the principal quantum number (n-1): ℓ =
0,1,2,3,4…,(n−1) 1.1.4 The Magnetic Quantum Number (mℓ) The
magnetic quantum number ml determines the number of orbitals and
their orientation within a subshell. Consequently, its value
depends on the orbital angular momentum quantum number ℓ. Given a
certain ℓ, mlml is an interval ranging from –l –l to +l +l, so it
can be zero, a negative integer, or a positive integer. m ℓ = −l,
(−l+1), (−l+2), …,−2, −1, 0,1,2,… (l–1), (l–2), +l 1.1.4.1 A Closer
Look at Shells and Subshells Principal Shells The value of the
principal quantum number n is the level of the principal electronic
shell (principal level). All orbitals that have the same n value
are in the same principal level. For example, all orbitals on the
second principal level have a principal quantum number of n=2. When
the value of n is higher, the number of principal electronic shells
is greater. This causes a greater distance between the farthest
electron and the nucleus. As a result, the size of the atom and its
atomic radius increases.
-
10 Because the atomic radius increases, the electrons are
farther from the nucleus. Thus it is easier for the atom to expel
an electron because the nucleus does not have as strong a pull on
it, and the ionization energy decreases. Subshells The number of
values of the orbital angular number l can also be used to identify
the number of subshells in a principal electron shell: • When n =
1, l= 0 (l takes on one value and thus there can only be one
subshell) • When n = 2, l= 0, 1 (l takes on two values and thus
there are two possible subshells) • When n = 3, l= 0, 1, 2 (l takes
on three values and thus there are three possible subshells) After
looking at the examples above, we see that the value of n is equal
to the number of subshells in a principal electronic shell: •
Principal shell with n = 1 has one subshell • Principal shell with
n = 2 has two subshells • Principal shell with n = 3 has three
subshells To identify what type of possible subshells n has, these
subshells have been assigned letter names. The value of l
determines the name of the subshell: Name of Subshell Value of ll s
subshell 0 p subshell 1 d subshell 2 f subshell 3 Therefore: •
Principal shell with n = 1 has one s subshell (l = 0) • Principal
shell with n = 2 has one s subshell and one p subshell (l = 0, 1) •
Principal shell with n = 3 has one s subshell, one p subshell, and
one d subshell (l = 0, 1, 2)
-
We can designate a principal quantum number, n, and a certain
subshell by combining the value of n and the name of the subshell
(which can be found usingto the third principal quantum number
(n=3) and the 1.1.5 Pauli's Exclusion PrincipleThe Pauli Exclusion
PrincipleIn other words, no electrons in an atom are permitted to
have an identical set ofnumbers. � The Pauli Exclusion
Principleelectron arrangements in atoms and molecules, and helping
to rationalize patterns in the periodic table. � In chemistry the
Pauli Exclusion Principleabout to discuss. � Wolfgang Pauli
received the 1945 Nobel Prize in Physics for his discovery as it
applied to electrons. � Later the Pauli Exclusion Principlewill
mention at the end of this page. 1.1.5.2 Four Quantum NumbersEvery
electron in an atom can be defined completely by four quantum
numbers:
� n: the principal quantum number� l: the orbital angular
momentum quantum number11 We can designate a principal quantum
number, n, and a certain subshell by combining the value of n and
the name of the subshell (which can be found using l). For example,
3p refers to the third principal quantum number (n=3) and the p
subshell (l=1). Pauli's Exclusion Principle Exclusion Principle
says that every electron must be in its own unique state. In other
words, no electrons in an atom are permitted to have an identical
set ofExclusion Principle sits at the heart of chemistry, helping
to explain the electron arrangements in atoms and molecules, and
helping to rationalize patterns in Exclusion Principle is applied
solely to electrons, which we are Wolfgang Pauli received the 1945
Nobel Prize in Physics for his discovery as it Exclusion Principle
was found to have a broader meaning, which we will mention at the
end of this page. Four Quantum Numbers Every electron in an atom
can be defined completely by four quantum numbers: n: the principal
quantum number the orbital angular momentum quantum number We can
designate a principal quantum number, n, and a certain subshell by
combining the l). For example, 3p refers says that every electron
must be in its own unique state. In other words, no electrons in an
atom are permitted to have an identical set of quantum sits at the
heart of chemistry, helping to explain the electron arrangements in
atoms and molecules, and helping to rationalize patterns in is
applied solely to electrons, which we are Wolfgang Pauli received
the 1945 Nobel Prize in Physics for his discovery as it roader
meaning, which we
-
� ml: the magnetic quantum number� ms: the spin quantum number
1.1.5.3 Example of the Pauli Exclusion Principle Consider argon's
electron configuration:The exclusion principle asserts that every
electron in an argon atom is in a unique sta12 he magnetic quantum
number : the spin quantum number Example of the Pauli Exclusion
Principle Consider argon's electron configuration: 1s2 2s2 2p6 3s2
3p6 The exclusion principle asserts that every electron in an argon
atom is in a unique sta The exclusion principle asserts that every
electron in an argon atom is in a unique state.
-
13 The 1s level can accommodate two electrons with identical n,
l, and ml quantum numbers. Argon's pair of electrons in the 1s
orbital satisfy the exclusion principle because they have opposite
spins, meaning they have different spin quantum numbers, ms. One
spin is +½, the other is -½. (Instead of saying +½ or -½ often the
electrons are said to be spin-up or spin-down .) The 2s level
electrons have a different principal quantum number to those in the
1s orbital. The pair of 2s electrons differ from each other because
they have opposite spins. The 2p level electrons have a different
orbital angular momentum number from those in the s orbitals, hence
the letter p rather than s. There are three p orbitals of equal
energy, the px, py and pz. These orbitals are different from one
another because they have different orientations in space. Each of
the px, py and pz orbitals can accommodate a pair of electrons with
opposite spins. The 3s level rises to a higher principal quantum
number; this orbital accommodates an electron pair with opposite
spins. The 3p level's description is similar to that for 2p, but
the principal quantum number is higher: 3p lies at a higher energy
than 2p. 1.1.6 Hund's Rule Hund's rule states that the lowest
energy electron configuration, the ground state, in any electron
subshell is the one with the greatest number of parallel electron
spins. Example 1 Consider the different ways in which a pair of
electrons might be arranged in p orbitals. Bearing in mind that we
need to satisfy the Pauli exclusion principle, the three possible
ways are shown below:
-
14 The middle option has the greatest number of parallel
electron spins and therefore has the lowest energy, i.e. Hund's
rule identifies the middle option as the electronic ground state.
It is the ground state because: • accommodating the electrons in
the same orbital (the leftmost option) increases the electrostatic
repulsion between the electrons because on average they are closer
together • electrons with opposite spins (the rightmost option)
tend to move closer to one another than if spins are parallel (spin
correlation). If electrons are closer, the electrostatic repulsion
between them increases Example 2 Here are some of the ways three
electrons could be placed into p orbitals. Options (b) and (e) have
the greatest number of parallel electron spins and therefore the
lowest energy. Options (b) and (e) have the same energy – they are
said to be degenerate orbitals. Example 3 Here are some of the ways
two electrons could be placed into d orbitals. Options (b),(c), and
(d) have the greatest number of parallel electron spins and
therefore the lowest energy. Options (b),(c), and (d) have the same
energy - they are said to be degenerate orbitals. 1.1.7 Aufbau
Principle Aufbau (German word) means ‘building up’ according to
Aufbau principle. “The vacant subshell having lowest energy is
filled first. When this subshell is filled completely, then the
filling of next sub-shell with higher energy starts.” The energy
level diagram is
-
The aufbau principle says that the arrangement of electrons in
an atom configuration - is best understood if it is built from the
ground up.atom's electron configuration, we begin at the lowest
eenergy sublevels until the required number of electrons are
present.We do so in accordance with theThe diagram below shows as
many electron energy levels as we need for most purposes in general
chemistry, although higher levels also exist. Energy levels are
filled from the bottom up. 15 The aufbau principle says that the
arrangement of electrons in an atom is best understood if it is
built from the ground up. When writing down an atom's electron
configuration, we begin at the lowest energy level and add
electrons to higher until the required number of electrons are
present. We do so in accordance with the Pauli Exclusion Principle
and Hund's RuleThe diagram below shows as many electron energy
levels as we need for most purposes in h higher levels also exist.
Energy levels are filled from the bottom The aufbau principle says
that the arrangement of electrons in an atom - the electron When
writing down an nergy level and add electrons to higher Hund's
Rule. The diagram below shows as many electron energy levels as we
need for most purposes in h higher levels also exist. Energy levels
are filled from the bottom
-
1.1.7.1 Element 1 - Hydrogen The simplest element, hydrogen, has
one electron.Hydrogen has a single electron. In its ground state,
hydrogen's electron occupies the 1s energy level. Note the energy
of 2s and 2p are equal for hydrogen This is not the case for any
other element's atoms.� The aufbau principle places hydrogen's
electron in the lowest available energy level, the 1s. � This is
hydrogen's ground state� Hydrogen is unique among the elements
electron. As a result of this, hydrogen has unique energy levels.�
The lack of electron-electron repulsion in hydrogen means its
sublevels are of equal energy: in hydrogen 2s and 2p are of equal
energy; likewise, 3s, 3p, and 3d are of equal energy, etc. 1.1.7.2
Element 2 - Helium Helium's two electrons fill the 1s orbital,
giving a 16 The simplest element, hydrogen, has one electron.
Hydrogen has a single electron. In its ground state, hydrogen's
electron occupies the 1s l. Note the energy of 2s and 2p are equal
for hydrogen - likewise 3s, 3p and 3d. This is not the case for any
other element's atoms. The aufbau principle places hydrogen's
electron in the lowest available energy level, ground state. It is
written 1s1. Hydrogen is unique among the elements - it is the only
one whose atoms have just one electron. As a result of this,
hydrogen has unique energy levels. electron repulsion in hydrogen
means its sublevels are of equal energy: in hydrogen 2s and 2p are
of equal energy; likewise, 3s, 3p, and 3d are of Helium's two
electrons fill the 1s orbital, giving a ground state electron
configuration of 1s Helium's ground state. Hydrogen has a single
electron. In its ground state, hydrogen's electron occupies the 1s
likewise 3s, 3p and 3d. The aufbau principle places hydrogen's
electron in the lowest available energy level, it is the only one
whose atoms have just one electron repulsion in hydrogen means its
sublevels are of equal energy: in hydrogen 2s and 2p are of equal
energy; likewise, 3s, 3p, and 3d are of ground state electron
configuration of 1s2.
-
� One electron is spin-up and the is other spinexclusion
principle. This is called spin pairing.� The quantum numbers of one
of helium's electrons are:n = 1, l = 0,� And the other electron's
quantum numbers are:n = 1, l = 0,� Helium's first energy shell
(with principal quantum number, n = 1) is full. This is a
particularly stable configuration 1.1.8 Madelung energy ordering
rule The order in which subshells are filled isthe Madelung rule
(after Erwin Madelungrule (after Charles Janet or Vsevolod
Klechkovskyspeaking, countries), or the diagonal ruleThe states
crossed by same red arrow have same n + l value.The direction of
the red arrow indicates the order of state filling.Orbitals with a
lower n + this context, n represents the number; the values ℓ = 0,
1, 2, 3 correspond to thesubshell ordering by this rule is 1s, 2s,
2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d6d, 7p, 8s, ... The rule is based on
the total number of nodes in the atomic orbital,related to the
energy.[3] In the case of equa 17 up and the is other spin-down in
accordance with the Pauli exclusion principle. This is called spin
pairing. of one of helium's electrons are: = 0, ml = 0, ms = +½ And
the other electron's quantum numbers are: = 0, ml = 0, ms = −½
Helium's first energy shell (with principal quantum number, n = 1)
is full. This is a particularly stable configuration, leading to
helium's lack of chemical reactivity.Madelung energy ordering rule
The order in which subshells are filled is given by the n + ℓ rule,
also known as Erwin Madelung), or the Janet ruleor the Klechkowsky
Vsevolod Klechkovsky in some, mostly French and Russiandiagonal
rule. sed by same red arrow have same n + l value. The direction of
the red arrow indicates the order of state filling. n + ℓ value are
filled before those with higher n + principal quantum number and ℓ
the azimuthal quantum = 0, 1, 2, 3 correspond to the s, p, d, and f
labels, respectively. The subshell ordering by this rule is 1s, 2s,
2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, The
rule is based on the total number of nodes in the atomic orbital, n
+ In the case of equal n + ℓ values, the orbital with a lowerdown
in accordance with the Pauli Helium's first energy shell (with
principal quantum number, n = 1) is full. This is a , leading to
helium's lack of chemical reactivity. , also known as Klechkowsky
in some, mostly French and Russian- The direction of the red arrow
indicates the order of state filling. n + ℓ values. In azimuthal
quantum labels, respectively. The , 5p, 6s, 4f, 5d, 6p, 7s, 5f, n +
ℓ, which is values, the orbital with a lower n value is
-
18 filled first. The fact that most of the ground state
configurations of neutral atoms fill orbitals following this n + ℓ,
n pattern was obtained experimentally, by reference to the
spectroscopic characteristics of the elements. The Madelung energy
ordering rule applies only to neutral atoms in their ground state.
There are 10 elements among the transition metalsand 9 elements
among the Lanthanoids and Actinoids for which the Madelung rule
predicts an electron configuration that differs from that
determined experimentally by one electron orbit (in the case of
Palladium and Thorium by two electron orbits). 1.1.9 Stability of
Half-Filled and Fully Filled Orbitals The exactly half-filled and
fully filled orbitals have greater stability than other
configurations. The reasons for their stability are symmetry and
exchange energy. (a) Symmetry The half-filled and fully-filled
orbitals are more symmetrical than any other configuration and
symmetry leads to greater stability. (b) Exchange Energy The
electrons present in the different orbitals of the same sub-shell
can exchange their positions. Each such exchange leads to the
decrease in energy known as Exchange Energy. Greater the number of
exchanges, greater the exchange energy and hence greater the
stability. As the number of exchanges that take place in the
half-filled and fully-filled orbitals is maximum, thus exchange
energy is maximum and hence maximum stability. 1.2. PERIODIC TABLE
1.2.1 The History of the Periodic Table The modern periodic table
has evolved through a long history of attempts by chemists to
arrange the elements according to their properties as an aid in
predicting chemical behavior. One of the first to suggest such an
arrangement was the German chemist Johannes Dobereiner (1780–1849),
who noticed that many of the known elements could be grouped in
triads, sets of three elements that have similar properties—for
example, chlorine, bromine, and iodine; or copper, silver, and
gold. Dobereiner proposed that all elements could be grouped in
such triads, but subsequent attempts to expand his concept were
unsuccessful. We
-
19 now know that portions of the periodic table—the d block in
particular—contain triads of elements with substantial
similarities. The middle three members of most of the other
columns, such as sulfur, selenium, and tellurium in group 16 or
aluminum, gallium, and indium in group 13, also have remarkably
similar chemistry. By the mid-19th century, the atomic masses of
many of the elements had been determined. The English chemist John
Newlands (1838–1898), hypothesizing that the chemistry of the
elements might be related to their masses, arranged the known
elements in order of increasing atomic mass and discovered that
every seventh element had similar properties ("The Arrangement of
the Elements into Octaves as Proposed by Newlands"). (The noble
gases were still unknown.) Newlands therefore suggested that the
elements could be classified into octaves, corresponding to the
horizontal rows in the main group elements. Unfortunately,
Newlands’s “law of octaves” did not seem to work for elements
heavier than calcium, and his idea was publicly ridiculed. At one
scientific meeting, Newlands was asked why he didn’t arrange the
elements in alphabetical order instead of by atomic mass, since
that would make just as much sense! Actually, Newlands was on the
right track—with only a few exceptions, atomic mass does increase
with atomic number, and similar properties occur every time a set
of ns2np6 subshells is filled. Despite the fact that Newlands’s
table had no logical place for the d-block elements, he was honored
for his idea by the Royal Society of London in 1887. The periodic
table achieved its modern form through the work of the German
chemist Julius Lothar Meyer (1830–1895) and the Russian chemist
Dimitri Mendeleev (1834–1907), both of whom focused on the
relationships between atomic mass and various physical and chemical
properties. In 1869, they independently proposed essentially
identical arrangements of the elements. Meyer aligned the elements
in his table according to periodic variations in simple atomic
properties, such as “atomic volume” (Figure 7.2 "Variation of
Atomic Volume with Atomic Number, Adapted from Meyer’s Plot of
1870"), which he obtained by dividing the atomic mass (molar mass)
in grams per mole by the density of the element in grams per cubic
centimeter. This property is equivalent to what is today defined as
molar volume (measured in cubic centimeters per mole): molar mass
(g/mol) / density (g/cm3) = molar volume (cm3/mol) "Variation of
Atomic Volume with Atomic Number, Adapted from Meyer’s Plot of
1870", the alkali metals have the highest molar volumes of the
solid elements. In Meyer’s
-
20 plot of atomic volume versus atomic mass, the nonmetals occur
on the rising portion of the graph, and metals occur at the peaks,
in the valleys, and on the downslopes. One group of elements that
is absent from Mendeleev’s table is the noble gases, all of which
were discovered more than 20 years later, between 1894 and 1898, by
Sir William Ramsay (1852–1916; Nobel Prize in Chemistry 1904).
Initially, Ramsay did not know where to place these elements in the
periodic table. Argon, the first to be discovered, had an atomic
mass of 40. This was greater than chlorine’s and comparable to that
of potassium, so Ramsay, using the same kind of reasoning as
Mendeleev, decided to place the noble gases between the halogens
and the alkali metals. 1.2.2 The Role of the Atomic Number in the
Periodic Table Despite its usefulness, Mendeleev’s periodic table
was based entirely on empirical observation supported by very
little understanding. It was not until 1913, when a young British
physicist, H. G. J. Moseley (1887–1915), while analyzing the
frequencies of x-rays emitted by the elements, discovered that the
underlying foundation of the order of the elements was by the
atomic number, not the atomic mass. Moseley hypothesized that the
placement of each element in his series corresponded to its atomic
number Z, which is the number of positive charges (protons) in its
nucleus. Argon, for example, although having an atomic mass greater
than that of potassium (39.9 amu versus 39.1 amu, respectively),
was placed before potassium in the periodic table. While analyzing
the frequencies of the emitted x-rays, Moseley noticed that the
atomic number of argon is 18, whereas that of potassium is 19,
which indicated that they were indeed placed correctly. Moseley
also noticed three gaps in his table of x-ray frequencies, so he
predicted the existence of three unknown elements: technetium (Z =
43), discovered in 1937; promethium (Z = 61), discovered in 1945;
and rhenium (Z = 75), discovered in 1925. 1.2.3 Atomic Volume
Definition The atomic volume is the volume one mole of an element
occupies at room temperature. Atomic volume is typically given in
cubic centimetres per mole - cc/mol. The atomic volume is a
calculated value using the atomic weight and the density using the
formula:
-
21 atomic volume = atomic weight/density Another way to
calculate atomic volume is to use the atomic or ionic radius of an
atom (depending on whether or not you are dealing with an ion).
This calculates is based on the idea of an atom as a sphere, which
isn't precisely accurate. However, it's a decent approximation. In
this case, the formula for the volume of a sphere is used: volume =
(4/3)(π)(r3) where r is the atomic radius For example, a hydrogen
atom has an atomic radius of 53 picometers. The volume of a
hydrogen atom would be: volume = (4/3)(π)(533) volume = 623000
cubic picometers (approximately) Units : cm3 Notes: The molar
volume is also known as the atomic volume. The standard SI units
are m3. Normally, however, molar volume is expressed in units of
cm3. To convert quoted values to m3, divide by 1000000. The molar
volume depends upon density, phase, allotrope, and temperature.
Values here are given, where possible, for the solid at 298 K.
1.2.4 Atomic Radius Atomic radius is the distance from the centre
of the nucleus to the outermost shell containing electrons. In
other words, it is the distance from the centre of the nucleus to
the point up to which the density of the electron cloud is
maximum.
-
1.2.4.1Types of Atomic Radii Atomic radii are divided into three
types:• Covalent radius• Van der Waals radius• Metallic
radiusTherefore, we will study these three types of understanding
of the subject. i) Covalent Radius Covalent radius is one half the
atoms of the same element in a molecule. Therefore, rtwo bonded
atoms). The internuclear distance between two bonded atoms is
called thelength. Therefore, ii) Van der Waals Radius It is one
half the distancesor two adjacent identical atoms belonging to two
neighbouring molecules of an element in the 22 Atomic radii are
divided into three types: Covalent radius Van der Waals radius
Metallic radius Therefore, we will study these three types of
radius because they are vital for the better Covalent radius is one
half the distances between the nuclei of two covalently bonded
atoms of the same element in a molecule. Therefore, r covalent = ½
(internuclear distance between two bonded atoms). The internuclear
distance between two bonded atoms is called ther covalent = ½( bond
length) distances between the nuclei of two identical non-bonded
isolated atoms or two adjacent identical atoms belonging to two
neighbouring molecules of an element in the radius because they are
vital for the better between the nuclei of two covalently bonded
clear distance between two bonded atoms). The internuclear distance
between two bonded atoms is called the bond onded isolated atoms or
two adjacent identical atoms belonging to two neighbouring
molecules of an element in the
-
23 solid state. The magnitude of the Van der Waals radius is
dependent on the packing of the atoms when the element is in the
solid state. For example, the internuclear distance between two
adjacent chlorine atoms of the two neighbouring molecules in the
solid state is 360 pm. Therefore, the Van der Waals radius of the
chlorine atom is 180 pm. iii) Metallic Radius A metal lattice or
crystal consists of positive kernels or metal ions arranged in a
definite pattern in a sea of mobile valence electrons. Each kernel
is simultaneously attracted by a number of mobile electrons and
each mobile electron is attracted by a number of metal ions. Force
of attraction between the mobile electrons and the positive kernels
is called the metallic bond. It is one half the internuclear
distances between the two adjacent metal ions in the metallic
lattice. In a metallic lattice, the valence electrons are mobile;
therefore, they are only weakly attracted by the metal ions or
kernels. In a covalent bond, a pair of electrons is strongly
attracted by the nuclei of two atoms. Thus, a metallic radius is
always longer than its covalent radius. For example, the metallic
radius of sodium is 186 pm whereas its covalent radius as
determined by its vapour which exists as Na2 is 154 pm. The
metallic radius of Potassium is 231 pm while its covalent radius is
203 pm. 1.2.4.2 Variation of Atomic Radii in the Periodic Table a)
Variation with period � The Covalent and Van der Waals radii
decrease with increase in atomic number as we move from left to
right in a period. The alkali metals at the extreme left of the
periodic table have the largest size in a period. The halogens at
the extreme right of the periodic table have the smallest size. The
atomic size of nitrogen is the smallest. After nitrogen, atomic
size increases for Oxygen and then decreases for fluorine. The size
of atoms of inert gases is larger than those of the preceding
halogens. � As we move from left to right in a period, nuclear
charge increases by 1 unit in each succeeding element while the
number of shells remains the same. This enhanced nuclear charge
pulls the electrons of all the shells closer to the nucleus. This
makes each individual shells smaller and smaller. This result in a
decrease of the atomic radius as we move from left to right in a
period.
-
24 � The atomic radius abruptly increases as we move from
halogens to the inert gas. This is because inert gases have
completely filled orbitals. Hence, the inter-electronic are
maximum. We express the atomic size in terms of Van der Waals
radius since they do not form covalent bonds. Van der Waals radius
is larger than the covalent radius. Therefore, the atomic size of
an inert gas in a period is much higher than that of preceding
halogen b) Variation within a Group The atomic radii of elements
increase with an increase in atomic number from top to bottom in a
group. As we move down the group, the principal quantum number
increases. A new energy shell is added at each succeeding element.
The valence electrons lie farther and farther away from the
nucleus. As a result, the attraction of the nucleus for the
electron decreases. Hence, the atomic radius increases. A Solved
Example for You Q: Why is the Van der Waals Radius always greater
than the Covalent Radius? Ans: The Van der Waals forces of
attraction are weak. Therefore, the internuclear distance in case
of atoms held by Van der Waal forces is much larger than those
between covalently bonded atoms. Since a covalent bond is formed by
the overlap of two half-filled atomic orbitals, a part of electron
cloud becomes common. Therefore, covalent radii are always smaller
than van der Waal radius. 1.2.5 Atomic Radius The atomic radius is
the distance from the atomic nucleus to the outermost stable
electron of a neutral atom. In practice, the value is obtained by
measuring the diameter of an atom and dividing it in half. But, it
gets trickier from there. The atomic radius is a term used to
describe the size of the atom, but there is no standard definition
for this value. Atomic radius may actually refer to the ionic
radius, as well as the covalent radius, metallic radius, or van der
Waals radius. Atomic radius is the distance from the centre of the
nucleus to the boundary of the electron cloud. The atomic radius is
in the Angstrom level. Although we define the atomic radius for a
single atom, it is hard to measure it for a single atom. Therefore,
normally the distance between the nuclei of two touching atoms is
taken and divided by two, to get the
-
25 atomic radius. Depending on the bonding between two atoms the
radius can be categorized as metallic radius, covalent radius, Van
der Waals radius, etc. Atomic radii increases as you go down in a
column in the periodic table, because new layers of electrons are
adding. From left to right in a row, atomic radii decrease (except
for noble gases). 1.2.6 Ionic Radius The ionic radius is half the
distance between two gas atoms that are just touching each other.
In a neutral atom, the atomic and ionic radiuses are the same, but
many elements exist as anions or cations. If the atom loses its
outermost electron (positively charged or cation), the ionic radius
is smaller than the atomic radius because the atom loses an
electron energy shell. If the atom gains an electron (negatively
charged or anion), usually the electron falls into an existing
energy shell so the size of the ionic radius and atomic radius are
comparable. Atoms can gain or lose electrons and form negative or
positive charged particles respectively. These particles are called
ions. When neutral atoms remove one or more electrons, it forms
positively charged cations. And when neutral atoms take up
electrons, they form negatively charged anions. Ionic radius is the
distance from the centre of a nucleus to the outer edge of the ion.
However, most of the ions do not exist individually. Either they
are bonded with another counter ion, or they have interactions with
other ions, atoms or molecules. Because of this, the ionic radius
of a single ion varies in different environments. Therefore, when
ionic radii are compared, the ions in similar environments should
be compared. There are trends in the ionic radii in the periodic
table. As we go down in a column, additional orbitals are added to
atoms; therefore, the respective ions also have additional
electrons. Thus, from top to bottom the ionic radii increase. When
we go from left to right across a row, there is a specific pattern
of ionic radii change. For example, in the 3rd row, sodium,
magnesium and aluminum make +1, +2 and +3 cations respectively. The
ionic radii of these three are gradually decreasing. As the numbers
of protons are higher than the number of electrons, nucleus tends
to pull the electrons more and more towards the centre, which
result in decreased ionic radii. However, the anions in the 3rd row
have considerably higher ionic radii compared to the cationic
radii. Starting from P3- the ionic radii decrease to S2- and to
Cl–. Reason for having a larger ionic radius in anions can be
explained by addition of electrons into outer orbitals.
-
What is the difference between Atomic Radius and Ionic Radius?�
Atomic radius is an indication of the size of an atom. Ionic radius
is an indication of the size of an ion. � A cation ionic radius is
smaller than that of the atomic radius. And anionic radius is
larger tha1.2.7 Ionisation Potential The ionization energy or
ionization potential is the energy necessary to remove an electron
from the neutral atom. It is a minimum for the alkali metals which
have a single electron outside a closed shell. Itfor the noble
gases which have closed shells.The ionization energy (IE) is
qualitatively defined as the amount ofremove the most loosely bound
electron, the valence electron, ofform a cation. The ionization
potential or otherwise known as the ionization energy, is the
measure of the amount of energy required to remove an electron from
a neutral atom or ground state. The first electron is removed from
potential. The basic equation for ionization energy is:The amount
of energy necessary changes each time an electron is let go, since
it becomes more difficult to remove electrons after one or more has
already been removed from the atom or molecule. 1.2.8 Electron
Affinity The electron affinity is defined as the energy change that
occurs when an atom gains an electron, releasing energy in the
process. Let's remember that an electron is negatively charged, so
when an atom gains an electron, it becomes aSince we are talking
about a chthere is an equation used to determine the electron
affinity:26 ifference between Atomic Radius and Ionic Radius?
Atomic radius is an indication of the size of an atom. Ionic radius
is an indication of the size of an ion. A cation ionic radius is
smaller than that of the atomic radius. And anionic radius is
larger than the atomic radius. The ionization energy or ionization
potential is the energy necessary to remove an electron from the
neutral atom. It is a minimum for the alkali metals which have a
single electron outside a closed shell. It generally increases
across a row on the periodic maximum for the noble gases which have
closed shells. (IE) is qualitatively defined as the amount of
energyremove the most loosely bound electron, the valence electron,
of an isolated gaseous atom to The ionization potential or
otherwise known as the ionization energy, is the measure of the
amount of energy required to remove an electron from a neutral atom
or ground state. The first electron is removed from the valence
shell and can be noted as the first ionization The basic equation
for ionization energy is: X → X+ + e- The amount of energy
necessary changes each time an electron is let go, since it becomes
more difficult to remove electrons after one or more has already
been removed from s defined as the energy change that occurs when
an atom gains an electron, releasing energy in the process. Let's
remember that an electron is negatively charged, so when an atom
gains an electron, it becomes a negative ion. Since we are talking
about a change in energy, when an electron is added to an atom,
there is an equation used to determine the electron affinity:
Atomic radius is an indication of the size of an atom. Ionic radius
is an A cation ionic radius is smaller than that of the atomic
radius. And anionic The ionization energy or ionization potential
is the energy necessary to remove an electron from the neutral
atom. It is a minimum for the alkali metals which have a single
generally increases across a row on the periodic maximum energy
required to an isolated gaseous atom to The ionization potential or
otherwise known as the ionization energy, is the measure of the
amount of energy required to remove an electron from a neutral atom
or ground state. the valence shell and can be noted as the first
ionization The amount of energy necessary changes each time an
electron is let go, since it becomes more difficult to remove
electrons after one or more has already been removed from s defined
as the energy change that occurs when an atom gains an electron,
releasing energy in the process. Let's remember that an electron is
negatively ange in energy, when an electron is added to an
atom,
-
27 This equation shows that electron affinity is equal to the
negative change in energy. Let's clarify the sign convention for
the energy change associated with the gain of an electron. Remember
that the definition of an electron affinity is the energy released,
so that means that the reaction is exothermic. If a reaction is
exothermic, the change in energy is negative. This means that the
electron affinity is positive. For example, the electron affinity
of chlorine has the negative sign, which shows us the energy that
is released to add one electron to an atom. The giving off of
energy is shown with a negative sign. Based on this sign
convention, this means that a higher electron affinity indicates
that an atom more easily accepts electrons. A lower electron
affinity indicates that an atom does not accept electrons as
easily. 1.2.9 Factors that Affect Electron Affinity There are two
factors that can affect electron affinity. These are atomic size
and nuclear charge. With regard to atomic size, let's think about a
magnet and a refrigerator. When a magnet is closer to the surface
of the refrigerator, you can clearly feel the pull of the
attraction between the magnet and the refrigerator. The farther the
magnet gets away from the fridge, the less you feel the attraction
or pull. When looking at a drawing of a smaller atom side by side
with a bigger atom, it can be seen that a smaller atom's outermost
shell is closer to the nucleus than that of a bigger atom. Just
like our magnet and refrigerator analogy, the electron will feel
more attraction to the nucleus if it is closer.
-
28 The smaller the atom is, the closer the outermost shell is;
therefore, it is a stronger attraction between the nucleus and the
incoming electron. That means the electron affinity is higher for
smaller atoms. When looking at the periodic table the atomic radius
increases from top to bottom, moving down a column; therefore, the
electron affinity increases from the bottom to the top of the
column. Nuclear charge also affects electron affinity. The nuclear
charge is also known as the atomic number, which is the same as the
number of protons. Protons are positive subatomic particles. The
more protons there are, the greater the attraction is to electrons.
Trends or patterns can be seen in the periodic table with regard to
electron affinity. We are looking at this in terms of across the
period, not down a group. From left to right, the nuclear charge
increases, resulting in a greater attraction to incoming electrons.
So, we can say that from left to right across a period, the
electron affinity increases upward.
-
29 1.2.10 Trends in the Periodic Table Atomic number : The
number of protons in an atom. Atomic mass : The number of protons +
neutrons. Valence electrons : The number of electrons in the
outermost orbit of an atom. (e.g. fluorine has 7 valence electrons)
Metallic character : Metals have characteristically weak force of
attraction between the nucleus and the outermost electrons. The
weaker this force, the greater the metallic character. Atomic
radius : The distance from the center of the atom to the outermost
electrons. Atomic radius indicates the size of an atom. Ionization
energy : The minimum amount of energy needed to remove the
outermost electron from an isolated atom of a gaseous sample of an
element. . Electron affinity : The tendency of an atom to attract
electrons. 1.2.11 Factors Affecting the Properties Some of the
properties of the elements are related to the force of attraction
between the nucleus and the outermost electrons. This force of
attraction is dependent on 2 factors. These are, in order of
importance: a) The distance between the nucleus and the outer
electrons b) The number of protons in the nucleus (the nuclear
charge) � Factors Affecting Periodicity Periodic Trends: Atomic
Radius: the distance from the center of the nucleus to the
outermost electron (how big an atom is) –Ionic Radius: the size of
positive or negative ions Ionization energy: the amount of energy
required to remove an electron from an atom or ion
Electronegativity: the attraction an atom has for electrons �
Factors Affecting Periodicity Factors Affecting Periodicity:
Effective Nuclear Charge: the total positive force of the nucleus
(the more protons, the higher the effective nuclear charge)
Electron Repulsion: result of like- charged electrons electrostatic
repulsion � Factors Affecting Periodicity Stable electron
configuration: atoms strive to form s 2 p 6 configurations (full
valence*, full octet*…*we’ll get to this) Number of Energy
-
30 Levels Electron Shielding Effect: inner electrons temporarily
“blocking” the nucleus’s electrostatic attraction for outer
electrons Proton : electron ratio. Importance of periodic table: It
serves as an easy way to predict the behavior of the elements as we
move through groups or periods. Let us discuss how these various
properties such as atomic radius, ionization energy, electron
affinity and electronegativity vary across the elements of groups
and periods. 1. Atomic radius: This is a measure of the radius of
the atom from the center of the nucleus to the outer most orbital.
Atomic radius reduces as we move from left to right in a period.
This is mainly due to inclusion of excessive protons which draw the
electrons very close to the nucleus. In case of groups, atomic
radius increases from top to bottom. This is due to increase in the
number of orbitals. With the increase in the number of orbitals,
the valence electrons remain far away from the nucleus. Hence,
nucleus cannot exert more attraction on the electrons, increasing
the atomic radius. 2. Ionization Energy: ionization energy refers
to the energy required to remove an electron from the valence
shell. Ionization energy and the atomic radius are inversely
related. Greater the atomic radius lesser will be ionization energy
and lesser the atomic radius, greater will the ionization energy.
With the increase in the atomic radius, the effect of nucleus on
valence electrons decreases. Hence, they can be removed very
easily. This trend is observed in case of groups i.e. ionization
energy decreases as move down the group. In case of periods, atomic
radius decreases as we towards right. Hence, protons exert more
force on the electrons making it difficult to remove an electron.
Thus, ionization increases from left to right in a period. 3.
Electronegativity: it refers to the ability of an atom to attract
electron so as to form an ionic bond. Electronegativity decreases
as we move top to bottom in a group due to lesser effect of nucleus
on the valence shells. Electronegativity increases as we move left
to right in a period due to increased effect of nucleus on the
valence electrons. 4. Electron affinity: electron affinity refers
to the change in energy when an electron is added to an atom to
form a negatively charged ion.
-
31 Electron affinity increases as we move from left to right in
a period. This is due to stronger attraction of the protons of the
nucleus. Thus, metals are said to have less electron affinity
compared to non-metals other than noble gases. Electron affinity
decreases as move down the group. This is due to lesser effect of
nucleus on the valence shell to attract the electrons. As can be
seen from the above discussion, it is very easy to describe the
variation of properties among closely related elements with the
help of periodic table. Solved problems 1. Periodic table helps to
predict which of the following properties of elements a. atomic
radius b. ionization energy c. a and b d. atomic number Answer: c
2. Ionization energy __________ as we move down the group a.
increases b. decreases c. cannot be predicted d. none Answer: b 3.
Electron affinity __________ as move towards right of a period a.
Increases b. Decreases c. Cannot be predicted d. None Answer: a 4.
Energy required to remove an electron is termed as a. Electron
affinity b. Ionization energy c. Electronegativity d. A or b
Answer: b
-
32 1.3 Inorganic Qualitative Analysis 1.3.1 Solubility Product
Principle and Qualitative Analysis Learning Objective Describe the
application of the solubility product principle in the qualitative
analysis of a solution. Key Points o In qualitative analysis, a
solution is treated with various reagents to test for the presence
of certain ions. o Solubility-product constants can be used to
devise methods for separating ions in a solution by selective
precipitation. o Cations are usually classified into six groups,
where each group has a common reagent which can be used to separate
them from the solution. Terms • Qualitative analysis Determination
of the identity of the chemical species in a sample. • Qualitative
inorganic analysis. A method of analytical chemistry which seeks to
find elemental composition of inorganic compounds. • Precipitation
a reaction that leads to the formation of a heavier solid in a
lighter liquid; the precipitate so formed at the bottom of the
container. 1.3.2 Qualitative Analysis Classical qualitative
inorganic analysis is a method of analytical chemistry that seeks
to find the elemental composition of inorganic compounds. It is
mainly focused on detecting ions in an aqueous solution. The
solution is treated with various reagents to test for reactions
characteristic of certain ions, which may cause color change, solid
forming, and other visible changes. 1.3.3 Solubility Product
Principle and Qualitative Analysis Solubility-product constants can
be used to devise methods for separating ions in a solution by
selective precipitation. Selective precipitation is used to form a
solid with one of the ions in solution without disturbing the other
ions. You can continue this method to effectively separate all of
the ions in a solution. The entire traditional qualitative-analysis
scheme is based on the use of these equilibrium constants to
determine the correct precipitating ions and the correct
strategy.
-
33 1.3.3.1 Chemical precipitation: Precipitation is the
formation of a solid in a solution or inside another solid during a
chemical reaction or by diffusion in a solid. Precipitation is used
in qualitative chemical analysis. The extent to which a solute will
dissolve in a solvent is called its solubility. The solubility of a
chemical is conventionally expressed as the maximum number of grams
of a chemical that will dissolve in l00g of solvent but conversion
to mol L-1 or g L-1 is simple and may be appropriate for some
applications (see below). Since solubility is temperature
dependent, is always quoted at a specific temperature. With a very
few exceptions, increasing the temperature of a solvent increases
the solubility of the solute. 1.3.3.2 Saturated solutions: For
practical purposes, a saturated solution is one in which no more
solute will dissolve. For example, the solubility of sodium
chloride in water is 35.6g per l00g at 25°Cand 39.1 g per 100g at
100°Cand both solutions are saturated solutions at their respective
temperatures. If the 100°C solution is cooled to 25°C, then 3.5 g
of NaCl crystals will precipitate from the solution, because the
solution at 25°C requires only 35.6 g of NaCI for saturation. This
process is the basis of purification of compounds by
recrystallization. 1.3.3.3 Solubility product: In dilute aqueous
solutions, it has been demonstrated experimentally for poorly
soluble ionic salts (solubilities less than 0.01 mol L-1) that the
mathematical product of the total molar concentrations of the
component ions is a constant at constant temperature. This product,
Ks is called the solubility product. Thus for a saturated solution
of a simple ionic compound AB in water, we have the dynamic
equilibrium: ABsolid ↔ A+(aq) + B−(aq) Where AB represents the
solid which has not dissolved, in equilibrium with its ions in the
aqueous saturated solution. Then: Ks = [A+] × [B−] For example,
silver chloride is a solid of solubility 0.000 15g per 100mL of
water in equilibrium with silver cations and chloride ions. Then:
Ks = [Ag+] × [Cl−] The solubility of AgCl is 0.0015gmL−1 (l0 ×
solubility per 100g, assuming that the density of water is
1.0gmL−1) and therefore the solubility of AgCl is 0.0015 ÷ 143.5 =
1.05 ×
-
34 10−5 mol L−1. Thus the saturated solution contains 1.05 ×
10−5 mol L−1 of Ag+ ions and 1.05 x 10−5 mol L−1 of Cl− ions and
the solubility product Ks is: Ks = (1.05 × 10−5) × (1.05 × 10−5) =
1.1 × 10−10 mol2 L−2 If the solid does not have a simple I: I ratio
of its ionic components, e.g. PbCb, then the solubility product is
given by: Ks = [Pb2+] × [Cl−]2 In general terms, the solubility
product for a compound MyNx, is given by: Ks = [M+]y × [N−]z The
practical effects of solubility products are demonstrated in the
detection of anions and cations by precipitation and in
quantitative gravimetric analysis (p. 139). For example, if dilute
aqueous solutions of silver nitrate (solubility 55.6 g per 100g of
water) and sodium chloride (solubility 35.6 g per 100g of water)
are mixed, an immediate white precipitate of AgCl is produced
because the solubility product of AgCI has been exceeded by the
numbers of Ag+ and Cl− ions in the solution, even though the ions
come from different 'molecules'. A saturated solution of AgCl is
formed and the excess AgCI precipitates out. The solubility product
of the other combinations of ions is not exceeded and thus sodium
and nitrate ions remain in solution. Even if the concentration of
Ag+ is extremely low, the solubility product for AgCl can be
exceeded by the addition of an excess of Cl− ions, since it is the
multiplication of these two concentrations which defines the
solubility product. Thus soluble chlorides can be used to detect
the presence of Ag+ ions and, conversely, soluble silver salts can
be used to detect Cl− ions, both quantitatively and qualitatively.
1.3.4 Principle of Elimination of interfering anions Removal of
interfering radicals before IIIrd group analysis What are the
interfering radicals? How do they interfere in systematic
separation of cationic radicals? Why is it necessary to remove them
before IIIrd gr analysis? Why don't they interfere in Ist or IInd
group analysis? Interfering radicals are oxalate, tartrate,
fluoride, borate and phosphate and they are anionic radicals. They
form complex with IIIrd gr group reagent ammonium chloride and
ammonium hydroxide. This leads to incomplete precipitation of IIIrd
group cations and causes
-
35 immature precipitation of IVth and Vth group cations in
alkaline medium. Let’s try to understand it. Oxalate, tartrate,
fluoride, borate, silicate and phosphate of the metals are soluble
in acidic medium. If you remember, for 1stand 2nd analysis medium
remain acidic (dilute HCl) that’s why they do not interfere then.
But for 3rd group analysis the medium becomes alkaline by group
reagents ammonium chloride and ammonium sulphide. Here interfering
radicals come into action and disturb the solubility product of
cations which causes their premature or incomplete precipitation.
In acidic medium these salts produce their corresponding acids like
oxalic acid, phosphoric acid, hydrofluoric acid, boric acid and
tartaric acid. For example, barium oxalate reacts with HCl and
produces oxalic acid. BaC2O4 + 2HCl ⟶ BaCl2 + H2C2O4 These
interfering acids are weak acids so they do not dissociate
completely and remain in solution in their unionised form.
Equilibrium is developed between dissociated and un-dissociated
acid. H2C2O4 ⇌ 2H+ + C2O42- Hydrochloric acid is a strong acid and
is ionised completely. HCl ⟶ H+ + Cl- Hydrogen ions acts as common
ion among them and higher concentration of H+ suppresses the
ionization of interfering acid. Therefore, ionic product of C2O42-
and Ba2+ doesn’t exceed the solubility product of barium oxalate
which is why Ba2+ remains in the solution as barium oxalate. That’s
how interfering radicals do not interfere as long as the medium
remains acidic enough. But when we make the medium alkaline by
adding 3rd group reagent ammonium hydroxide NH4OH, OH- ions combine
with H+and neutralise them. This decreases the concentration of H+
ions which shifts the equilibrium of dissociation of interfering
acid forward and increases the concentration of C2O42- . Thus the
ionic product of C2O42- and Ba2+ exceeds the solubility product of
barium oxalate and Ba2+ gets precipitated in the 3rd group, which
actually belongs to the 4th group.
-
36 One or more interfering radicals can be present in the
solution. They have to be removed in the following order: first we
remove oxalate and tartrate, then borate and fluoride, then
silicate and in the last phosphate. i) Procedure for the removal of
oxalate and tartrate: Oxalate and tartrate of metals are soluble in
acid and they decompose on heating. Take the filtrate of 2nd group
and boil off H2S gas from it. Add 4-5ml concentrated nitric acid
HNO3 and heat it till it is almost dry. Repeat this treatment for
2-3 times. (COO)22- + H+ ⟶ (COOH)2 (COOH)2 ⟶ HCOOH + CO2↑ HCOOH ⟶
CO↑ + H2O↑ Tartrate and tartaric acid decomposes in a complex
manner; charring takes place on heating and a smell of burnt sugar
develops. Extract the with dilute HCl and filter. Use this filtrate
for analysis of 3rd group or use for removal of other interfering
radicals. ii) Procedure for the removal of borate and fluoride:
Take the filtrate and evaporate it to dryness. Add concentrated HCl
and again evaporate to dryness. F- + H+ ⟶ HF CaF2 + 2HCl ⟶ CaCl2 +
2HF On heating with HCl fluoride forms hydrofluoric acid and Borate
forms orthoboric acid which evaporate on heating. BO33- + 3H+ ⟶
H3BO3 Na3BO3 + 3HCl ⟶ 3NaCl + H3BO3 Extract the residue with dilute
HCl and filter. Use this filtrate for analysis of 3rd group or use
for removal of other interfering radicals. If fluoride is absent
and borate is present then residue use a mixture of 5ml ethyl
alcohol and 10ml conc. HCl and evaporate to dryness. BO33- + 3H+ ⟶
H3BO3 H3BO3 + 3C2H5OH ⟶ (C2H5O)3B↑ + H2O iii) Procedure for the
removal of silicate: Evaporate the filtrate of 2nd group or residue
obtained from removal of interfering radicals with concentrated HCl
to dryness. Repeat this treatment for 3-4 times.
-
37 SiO32- + 2H+ ⟶ H2SiO3 ↓ H2SiO3 ↓ ⟶ SiO2 ↓ + H2O On heating
with HCl silicate converts to metasilicic acid (H2SiO3) which is
converted into white insoluble powder silica (SiO2) on repetitive
heating with concentrated HCl. iv) Test for phosphate HPO42-: test
0.5ml of the filtrate with 1ml ammonium molybdate reagent and a few
drops of concentrated HNO3, and warm gently, yellow precipitate
indicates the presence of phosphate. Its composition is not known
exactly. v) Procedure for the removal of phosphate: Ferric chloride
is generally used for the removal of phosphate. Fe(III) combines
with phosphate and removes all phosphate as insoluble FePO4.
Fe(III) is also a member of 3rd group so first we have to test its
presence in the filtrate of 2nd group then we can proceed for the
removal of phosphate. HPO42- + Fe3+ ⟶ FePO4 ↓ + H+ Test for Fe: To
the filtrate of 2nd group add ammonium chloride NH4Cl and a slight
excess of ammonia NH3 solution. If precipitate appears, it
indicates the presence of 3rd group. It may contain hydroxides
Fe(OH)3, Cr(OH)3, Al(OH)3, MnO2.xH2O, traces of CaF2 and phosphates
of Mg and IIIA, IIIB and IV group metals. Dissolve the precipitate
in minimum volume of 2M HCl. Take 0.5ml solution and add potassium
hexacyanoferrate (II) K4[Fe(CN)6] solution. If iron is present, you
will get prussian blue coloured precipitate of iron(III)
hexacyanoferrate. 4Fe3+ + 3[Fe(CN)6]4- ⟶ Fe4[Fe(CN)6]3 Removal of
phosphate: To the main solution add 2M ammonia NH3 solution drop
wise, with stirring, until a faint permanent precipitate is just
obtained. Then add 2-3ml 9M acetic acid and 5ml 6M ammonium acetate
solution. Discard any precipitate if obtained at this stage. If the
solution is red or brownish red, sufficient iron Fe(III) is present
here to combine with phosphate. If the solution is not red or
brownish red in colour then add ferric chloride FeCl3 solution drop
wise with stirring, until the solution gets a deep brownish red
coloured. Dilute the solution to about 150ml with hot water, boil
gently for 1-2min, filter hot and wash
-
38 the residue with a little boiling water. Residue may contain
phosphate of Fe, Al and Cr. Keep the filtrate for test of IIIB
group. Rinse the residue in porcelain dish with 10ml cold water,
add 1-1.5g sodium peroxoborate and boil gently until the evolution
of O2 ceases (2-3min). Filter and wash with hot water. Reject the
residue to remove phosphate in the form of FePO4. Keep the filtrate
and test for IIIA group. To test the presence of interfering
radicals you need to prepare sodium carbonate extract and then test
them separately. Scheme for the test of anionic radicals is not as
systematic as cationic radicals. We will study them in coming
posts. In the next post we will discuss the analysis of IIIA group
cations. 1.3.5 Common-ion effect The common-ion effect states that
in a chemical solution in which several species reversibly
associate with each other by an equilibrium process, increasing the
the concentration of any one of its dissociated components by
adding another chemical that also contains it will cause an
increased amount of association. This result is a consequence of Le
Chatelier's principle for the equilibrium reaction of the
association/dissociation. The effect is commonly seen as an effect
on the solubility of salts and other weak electrolytes. Adding an
additional amount of one of the ions of the salt generally leads to
increased precipitation of the salt, which reduces the
concentration of both ions of the salt until the solubility
equilibrium is reached. The effect is based on the fact that both
the original salt and the other added chemical have one ion in
common with each other. 1.3.6 Uncommon – ion effect Sometimes
adding an ion other than the ones that are part of the precipitated
salt itself can increase the solubility of the salt. This "salting
in" is called the "uncommon-ion effect" (also "salt effect" or the
"diverse-ion effect"). It occurs because as the total ion
concentration increases, inter-ion attraction within the solution
can become an important factor. This alternate equilibrium makes
the ions less available for the precipitation reaction. This is
also called odd ion effect. 1.3.7 Solubility effects A practical
example used vary widely in areas drawing drinking water from chalk
or limestone aquifers is the addition of sodium carbonate to the
raw water to
-
39 reduce the hardness of the water. In the water treatment
process, highly soluble sodium carbonate salt is added to
precipitate out sparingly soluble calcium carbonate. The very pure
and finely divided precipitate of calcium carbonate that is
generated is a valuable by-product used in the manufacture of
toothpaste. The salting-out process used in the manufacture of
soaps benefits from the common-ion effect. Soaps are sodium salts
of fatty acids. Addition of sodium chloride reduces the solubility
of the soap salts. The soaps precipitate due to a combination of
common-ion effect and increased ionic strength. Sea, brackish and
other waters that contain appreciable amount of sodium ions (Na+)
interfere with the normal behavior of soap because of common-ion
effect. In the presence of excess Na+, the solubility of soap salts
is reduced, making the soap less effective. Solubility refers to
the amount of material that is able to be dissolved in a particular
solvent. For example, table salt (NaCl) placed in water eventually
dissolves. However, if more table salt is continuously added, the
solution will reach a point at which no more can be dissolved; in
other words, the solution is saturated, and the table salt has
effectively reached its solubility limit. 1.3.8 Buffering effect A
buffer solution contains an acid and its conjugate base or a base
and its conjugate acid.[2] Addition of the conjugate ion will
result in a change of pH of the buffer solution. For example, if
both sodium acetate and acetic acid are dissolved in the same
solution they both dissociate and ionize to produce acetate ions.
Sodium acetate is a strong electrolyte, so it dissociates
completely in solution. Acetic acid is a weak acid, so it only
ionizes slightly. According to Le Chatelier's principle, the
addition of acetate ions from sodium acetate will suppress the
ionization of acetic acid and shift its equilibrium to the left.
Thus the percent dissociation of the acetic acid will decrease, and
the pH of the solution will increase. The ionization of an acid or
a base is limited by the presence of its conjugate base or acid.
NaCH3CO2 (s) → Na+ (aq) + CH3CO2− (aq) CH3CO2H (aq) ⇌ H+(aq) +
CH3CO2−(aq) This will decrease the hydronium concentration, and
thus the common-ion solution will be less acidic than a solution
containing only acetic acid.
-
40 1.3.9 Complexation reactions including spot test in
qualitative analysis Spot Tests General Procedures Use clean
glassware and deionized water to avoid contaminating your sample.
Perform tests on “known” solutions and your unknown solution
simultaneously. Keep in mind that your unknown is diluted compared
and may be colored by the presence of several ions. Therefore,
positive results might have slightly different appearances between
the known and your unknown. Some spot test reactions occur slowly.
If no reaction is apparent after the addition of reagent, heat the
reaction mixture in a hot water bath according to instructions
given. Stir well after addition of reagent using a clean stirring
rod or disposable transfer pipet. Perform tests in the order given
as described in the Laboratory Manual a) Spot Tests Test 1:
Carbonate Ion You will first test the “known” solution by placing 5
drops of the 1 M Na2CO3 solution in a micro test tube (or spot
plate) and add 5 drops of 6 M HCl. The formation of bubbles
indicates release of carbon dioxide gas. Repeat the test using 5
drops of your unknown and adding 5 drops of 6 M HCl (do NOT add the
1 M Na2CO3 solution as this is the “known” solution). b) Spot Tests
Test 2: Thiocyanate Ion This basis of the thiocyanate spot test is
to identify the thiocyanate ion, SCN- , by its reaction with
iron(III) to form a blood-red colored complex. Add 5 drops of 6 M
HCl to 5 drops of 0.5 M KSCN (the “known”) on the spot plate. Add 1
drop of 0.1 M Fe(NO3)3 to the solution in the spot plate. The
solution will turn red if SCN- is present. Repeat the test on your
unknown solution without adding the “known” 0.5 M KSCN solution. c)
Spot Tests Test 3: Sulfate Ion Put 5 drops of the “known” 0.5 M
Na2SO4 solution in a well of a spot plate. Add 5 drops of 6 M HCl
to the plate. Now add 1 drop of 1 M BaCl2 to the plate. A
precipitate of white BaSO4 indicates the presence of sulfate
ion.Repeat the test on your unknown without using the 0.5 M Na2SO4
“known” solution.
-
41 d) Spot Tests Test 4: Phosphate Ion Place 5 drops of the
“known” 0.5 M Na2HPO4 (sodium monohydrogen phosphate) solution in a
medium test tube. Add 1 mL of 6 M HNO3. Now add 1 mL of Ammonium
molybdate solution, (NH4)2MoO4 to the test tube. The presence of
phosphate ions is confirmed by the formation of a yellow
precipitate.This precipitate is often slow to form and heating in
the hot water bath may be required to initiate precipitation.Repeat
the test on your unknown. e) Spot Tests Test 5: Chromate Ion
Solutions containing chromate ions are yellow when neutral or basic
and orange when acidic. The orange color is due to the formation of
the dichromate ion (Cr2O72-).Add 5 drops of “known” 0.5 M K2CrO4 to
a spot plate. Add 5 drops of 3% H2O2. Now add 10 drops of 6 M HNO3
to the plate. The formation of a blue-green color, which appears
but rapidly disappears, indicates the presence of the chromate ion.
Repeat this test on your unknown without adding the 0.5 M K2CrO4
“known” solution. f) Spot Tests Test 6: Chloride Ion Place 5 drops
of “known” 0.5 M NaCl on a spot plate. Add 5 drops of 6 M HNO3. Now
add 2 drops of 0.1 M AgNO3 to the plate. The presence of chloride
ion is indicated by the formation of a white precipitate of AgCl.
If Test 2 indicated that your unknown contained SCN-, follow the
lab manual procedure for removing the SCN- ion before you test your
unknown for Cl- because SCN- also forms a white precipitate with
Ag+. Repeat this test on your unknown solution. Remember not to add
the “known” NaCl solution. g) Spot Tests Test 7: Acetate Ion Place
5 drops of “known” 0.5 M NaC2H3O2 (sodium acetate) on a spot plate.
(The acetate ion is often abbreviated as OAc-, or as CH3COO-). Add
5 drops of 3 M H2SO4 to the spot plate. The presence of acetate ion
is indicated by the smell of acetic acid (a vinegar odor) coming
from the plate. If no odor is immediately present, repeat the test
in a micro test tube, warming it in a hot water bath for 20 seconds
and check the odor again.Repeat this test on your unknown solution.
Do not add the “known” NaC2H3O2 solution.
-
42 h) Spot Tests Test 8: Ammonium Ion This test also relies on
your sense of smell to detect the presence of a substance. Add 5
drops of “known” 0.5 M NH4Cl to a spot plate or a micro test tube.
Add 5 drops of 6 M NaOH to the plate and check for the presence of
an ammonia odor by gentle wafting. If no odor is immediately
apparent, warm the test tube and check again.As an additional test
for NH4+, place a drop of the unknown on a spot plate and add a
drop of Nessler’s Reagent. A reddish-brown precipitate indicates
the presence of the NH4+ ion (see lower photo). Repeat the test on
your unknown. i) Spot Tests Interfering Ions If you suspect that an
ion is interfering with a particular test, make up a “known”
solution that contains the ions you think you have in our unknown
and repeat the spot tests which gave you trouble. Test 3 is most
likely to have interference problems from the Ba2+ ion which also
forms precipitates with the chromium and phosphate ions. Test 6 is
most likely to have interference problems from the Ag+ ion which
forms precipitates with the phosphate, sulfate, carbonate,
chromate, and thiocyanate ions. Refer to the Laboratory Manual for
specific instructions of how to remove these interfering ions.
1.3.10 Reactions involved in separation and identification of
cations and anions in the analysis Qualitative Analysis of Cation
and Anion Qualitative analysis is used to identify and separate
cations and anions in a sample substance. Unlike quantitative
analysis, which seeks to determine the quantity or amount of
sample, qualitative analysis is a descriptive form of analysis. In
an educational setting, the concentrations of the ions to be
identified are approximately 0.01 M in an aqueous solution. The
'semimicro' level of qualitative analysis employs methods used to
detect 1-2 mg of an ion in 5 mL of solution. While there are
qualitative analysis methods used to identify covalent molecules,
most covalent compounds can be identified and distinguished from
each other using physical properties, such as index of refraction
and melting point.
-
43 1.3.11 Sample Qualitative Analysis Protocol First, ions are
removed in groups from the initial aqueous solution. After each
group has been separated, then testing is conducted for the
individual ions in each group. Here is a common grouping of
cations: Group I: Ag+, Hg22+, Pb2+ Precipitated in 1 M HCl Group
II: Bi3+, Cd2+, Cu2+, Hg2+, (Pb2+), Sb3+ and Sb5+, Sn2+ and Sn4+
Precipitated in 0.1 M H2S solution at pH 0.5 Group III: Al3+,
(Cd2+), Co2+, Cr3+, Fe2+ and Fe3+, Mn2+, Ni2+, Zn2+ Precipitated in
0.1 M H2S solution at pH 9 Group IV: Ba2+, Ca2+, K+, Mg2+, Na+,
NH4+ Ba2+, Ca2+, and Mg2+ are precipitated in 0.2 M (NH4)2CO3
solution at pH 10; the other ions are soluble Many reagents are
used in the qualitative analysis, but only a few are involved in
nearly every group procedure. The four most commonly used reagents
are 6M HCl, 6M HNO3, 6M NaOH, 6M NH3. Understanding the uses of the
reagents is helpful when planning an analysis. It is the intention
of this unit to study the separation and identification of 21 of
the most common cations and 8 of the most common anions. Common
Cations: Ag1+, Hg22+, Pb2+, Hg2+, Bi3+, Cu2+, Cd2+, Al3+, Cr3+,