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+,