lectures part 1

7/22/2019 Lectures Part 1

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Vinnitsa National Pirogov Memorial Medical University

Biological and General Chemistry Department

Medical chemistry course

SYSTEMATIC COURSE

Lecture material

Module 1. Acid-base equilibrium and chemistry of complexes

in biological liquids

For practical lessons of medical chemistry for foreign students

Vinnitsa 2010


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Systematic course is approved by academic council of National Pirogov Memorial Medical

University,Vinnitsa (minutes 5 from 2.03.2011)

Authors:

Assistant Professor Smirnova O.V.Assistant Professor Chervyak M.M.

Assist. Shunkov V.S.

Reviewer:

Azarov O.S.- Candidate of chemistry science, assistant professor

Department of Biological and General Chemistry VNMU

Marchak T.V.- Candidate of chemistry science, assistant professor

Department of Physiological Agriculture and Live Stock Breeding

and Chemistry VNAU

Shitova T.V. Senior lecturer

Department of Russian and Ukrainian languages

Head of English language courses VNMU

Printing group VNMU:

Text editor Shunkov V.S.

Computer editor Shunkov V.S.

Secretary Koroleva N.D.


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CONTENT

1. Periodic system of D.I. Mendeleev. Electron-atomic

structure of elements and ions.

4

2. Biogenic s-elements, chemical properties, biological

role, application in medicine.

11

3. Biogenic-elements, chemical properties, biological


20

4. Biogenic d- elements, chemical properties, biological


33

5. The formation of complexes in biological systems. 38

6. Preparation of the solutions and calculation of its

concentrations.

46

7. The basic concepts of volumetric analysis.

Neutralization method. Base Standardization.

56

8. Neutralization method. Acid Standardization. 64

9. Acid-base equilibrium in human body. pH of

biological liquids.

67

10. Buffer systems, classification and mechanism. 71

11. Buffer capacity. The role of buffer solutions in

biological systems.

74

12. Colligative properties. Osmosis. 80


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TOPIC 1: Periodic system of D.I. Mendeleev. Electron-atomic structure of

elements and ions.

The periodic table of the chemical elements (also periodic table of the elements or just theperiodic table) is a tabular display of the chemical elements. Although precursors to this table exist,

its invention is generally credited to Russian chemist Dmitri Mendeleevin 1869, who intended thetable to illustrate recurring ("periodic") trends in the properties of the elements. The layout of thetable has been refined and extended over time, as new elements have been discovered, and newtheoretical models have been developed to explain chemical behavior.

The periodic table is now ubiquitous within the academic discipline of chemistry, providing auseful framework to classify, systematize, and compare all of the many different forms of chemical

behavior. The table has found many applications in chemistry, physics, biology, and engineering,especially chemical engineering. The current standard table contains 118 elementsto date

Classification:

Groups

A group or family is a vertical column in the periodic table. Groups are considered the mostimportant method of classifying the elements. In some groups, the elements have very similar

properties and exhibit a clear trend in properties down the group. These groups tend to be giventrivial (unsystematic) names, e.g., the alkali metals, alkaline earth metals, halogens, pnictogens,chalcogens, and noble gases. Some other groups in the periodic table display fewer similaritiesand/or vertical trends (for example Group 14), and these have no trivial names and are referred tosimply by their group numbers.

Periods

A period is a horizontal row in the periodic table. Although groups are the most common way of

classifying elements, there are some regions of the periodic table where the horizontal trends andsimilarities in properties are more significant than vertical group trends. This can be true in the d-block (or "transition metals"), and especially for the f-block, where the lanthanides and actinidesform two substantial horizontal series of elements.

Blocks

Because of the importance of the outermost shell, the different regions of the periodic table aresometimes referred to as periodic table blocks, named according to the subshell in which the "last"electron resides. The s-block comprises the first two groups (alkali metals and alkaline earth metals)as well as hydrogen and helium. The p-block comprises the last six groups (groups 13 through 18)and contains, among others, all of the semimetals. The d-block comprises groups 3 through 12 and

contains all of the transition metals. The f-block, usually offset below the rest of the periodic table,comprises the rare earth metals.The chemical elements are also grouped together in other ways. Some of these groupings are

often illustrated on the periodic table, such as transition metals, poor metals, and metalloids. Otherinformal groupings exist, such as the platinum group and the noble metals. (Table 1.1)


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Main-group elements

s-block

1A

Transition elements

p-block 8A1 2A 3A 4A 5A 6A 7A

2 d-block3 3B 4

B

5B 6B 7B 8B 1B 2B

4

5

6

7

f-block 45

Table 1.1

Russian chemistry professor Dmitri Ivanovich Mendeleev and JuliusLothar Meyer independently published their periodic tables in 1869 and1870, respectively. They both constructed their tables in a similar manner:

by listing the elements in a row or column in order of atomic weight andstarting a new row or column when the characteristics of the elements

began to repeat. The success of Mendeleev's table came from twodecisions he made: The first was to leave gaps in the table when it seemedthat the corresponding element had not yet been discovered. Mendeleevwas not the first chemist to do so, but he went a step further by using thetrends in his periodic table to predict the properties of those missingelements, such as gallium and germanium. The second decision was tooccasionally ignore the order suggested by the atomic weights and switch adjacent elements, suchas cobalt and nickel, to better classify them into chemical families. With the development oftheories of atomic structure, it became apparent that Mendeleev had inadvertently listed the

elements in order of increasing atomic number.

Periodicity of chemical properties :

The main value of the periodic table is the ability to predict the chemical properties of anelement based on its location on the table. It should be noted that the properties vary differentlywhen moving vertically along the columns of the table than when moving horizontally along therows.

Trends of groups

Modern quantum mechanical theories of atomic structure explain group trends by proposing thatelements within the same group have the same electron configurations in their valence shell, which

is the most important factor in accounting for their similar properties.Elements in the same group also show patterns in their atomic radius, ionization energy, and

electronegativity. From top to bottom in a group, the atomic radii of the elements increase. Sincethere are more filled energy levels, valence electrons are found farther from the nucleus.

- s-elements- p-elements- d-elements- f-elements


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From the top, each successive element has a lower ionization energy because it is easier toremove an electron since the atoms are less tightly bound. Similarly, a group will also see a top to

bottom decrease in electronegativity due to an increasing distance between valence electrons andthe nucleus.

Trends of periods

Elements in the same period show trends in atomic radius, ionization energy, electron affinity,

and electronegativity. Moving left to right across a period, atomic radius usually decreases. Thisoccurs because each successive element has an added proton and electron which causes the electronto be drawn closer to the nucleus. This decrease in atomic radius also causes the ionization energyto increase when moving from left to right across a period. The more tightly bound an element is,the more energy is required to remove an electron. Similarly, electronegativity will increase in thesame manner as ionization energy because of the amount of pull that is exerted on the electrons bythe nucleus. Electron affinity also shows a slight trend across a period. Metals (left side of a period)generally have a lower electron affinity than nonmetals (right side of a period) with the exception ofthe noble gases.

Electronic structure of atoms and ions.

To understand why pure substances have particular compositions and properties we need toknow about the "electronic structure" of the atoms (i.e. the way the electrons are arranged about thenucleus of the atoms of different elements). Then we can rationalize the ratio in which atomscombine and whether they are molecular, polymeric or ionic, [e.g. why gaseous nitrogen consists ofdiscrete N2 molecules; methane, ammonia, water and hydrogen fluoride consist of discretemolecules of CH4, NH3, H2O, HF respectively; why sodium chloride consists of Na+ and Cl ions;why metals exist in nature mainly as cations, Mx+; why free-radicals are reactive.

Early last century a model of the atom (Bohr model) in which electrons circulated around thenucleus in orbits, just as planets do around the sun, was developed. This model proved inadequateto explain many phenomena and was replaced in the 1920's, when it waspostulated that the

behaviour of electrons in atoms could be described by mathematicalequations similar to those used

to describe the motion of standing waves in a string. Fromthis model the electrons can be visualisedas electron clouds of various shapes with thenucleus of the atom at their centre.

Electronic structure of atoms: The arrangement of electrons around the nucleus of the atom.The properties of atoms can be understood in terms ofQuantum Theory, which involves theHeisenberg Uncertainty Principleand theSchrdinger Wave Equation.

Quantum Theory: A theory that states that the energy of an object can only change by discretesteps. A change involves a packet of energy called a quantum.

Heisenberg Uncertainty Principle: The position and momentum of a particle cannot both be

known simultaneously. This implies that in an atom the position and momentum of an electroncannot both be known simultaneously. (Thus a model of an atom containing electrons in fixedorbits around the nucleus is untenable.)

Schrdinger Wave Equation: A mathematical expression ascribing wave-like properties tomatter. When applied to atoms it describes the properties of electrons in atoms. This equation givesrise to the concepts ofenergy levels, atomic orbitals and quantum numbers.

Electronic energy levels: Allowed energies of electrons in atoms.

Atomic orbital: A mathematical expression from theSchrdinger Wave Equationfrom which,

for each energy level, the probability of finding the electron at different positions from the nucleuscan be calculated. The atomic orbital can be depicted as an "electroncloud" with the nucleus at thecentre, the denser the cloud the greater the probability of the electron being there.

Only two electrons can occupy the same orbital.


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Quantum numbers: Numbers which label the orbital and spin of an electron.

Electron pair: Two electrons in the same orbital. They must have opposite spins.

Spin of an electron: The intrinsic angular momentum of an electron. Occurs in only twosenses denoted () and ()).

Electron shells: The electrons in an atom exist in shells, each shell being made up of atomicorbitals or subshells.

Principal quantum number: Symboln, an integer, 1,2 3... which defines the shell. The smallern is, the lower the energy of the electron (more energy required to remove the electron from theatom), and the closer on average it is to the nucleus. First character in designation of an orbital.

Azimuthal quantum number: Symbol l, defines the subshell or kind of orbital, and can havethe values 0,1,...,n-1. An orbital with l= 0 is called ans orbital; with l= 1 is called a p orbital; withl= 2 is called a dorbital; with l= 3 is called an forbital. Second character in designation of anorbital.

Magnetic quantum number: Symbol m1, specifies the particular orbital of a subshell and canhave values -l, -l+1,...0,...,l-1,l.

Spin quantum number: Symbolms, specifies the spin of an electron and can have values of+ () or - ().

Occupancy of shells: The first shell, n = 1, can hold 2 electrons in one orbital, labelled 1s. (l= 0for ans orbital)

The second shell, n = 2, can hold 8 electrons in four orbitals, one labelled 2sand three labelled2p. (l= 1 for ap orbital).

The third shell can hold 18 electrons in nine orbitals, one 3s, three 3p and five 3d. (l= 2 for adorbital)

The fourth shell,n = 4, can hold 32 electrons in 16 orbitals, one 4s, three 4p, five 4dand seven4f.(l= 3 for anforbital)

Within a shell the energy levels of the orbitals (subshells) iss


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Quantum numbers

n, l

Spectroscopic

notation

or subshell ( n, l )

Maximum number

Of electrons allowed

In the subshell

= 2 (2l + 1)

6,2 6d 105,3 5f 14

7,0 7s 26,1 6p 65,2 5d 10

4,3 4f 146,0 6s 25,1 5p 64,2 4d 105,0 5s 24,1 4p 63,2 3d 10

4,0 4s 23,1 3p 63,0 3s 2

2,1 2p 62,0 2s 21,0 1s 2

Table 1.2

Atomic number

The atomic number of an element, Z, is equal to the number of protons that defines the element.For example, all carbon atoms contain 6 protons in their nucleus; so the atomic number "Z" of

carbon is 6. Carbon atoms may have different numbers of neutrons; atoms of the same element havingdifferent numbers of neutrons are known as isotopes of the element.

Atomic mass

The mass number of an element, A, is the number of nucleons (protons and neutrons) in the atomicnucleus. Different isotopes of a given element are distinguished by their mass numbers, which areconventionally written as a super-index on the left hand side of the atomic symbol (e.g., 238U).

In general, it differs slightly from the mass number as the mass of the protons and neutrons is notexactly 1 u, the electrons also contribute slightly to the atomic mass, and because of the nuclear binding

energy. For example, the mass of 19F is 18.9984032 u. The only exception to the atomic mass of anisotope not being a natural number is 12C, which has a mass of exactly 12, because u is defined as 1/12thof the mass of a free carbon-12 atom.

Basic Atomic Structure

The number of protons in an atom or ion determines what element it is.(Table 1.3)For example, if a particle has 6 protons in it, it must be carbon.


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(Subatomic Particles)

Subatomic Particle Mass Charge

Proton ~ 1 a.m.u. +1Electron ~ 0 a.m.u. -1Neutron ~ 1 a.m.u. 0

Table 1.3

Basic Atomic Structure

The atomic number of an atom or an ion is equal to its number of protons.Atomic Number = Number of Protons

Example:

If you look up potassium (K) in the periodic table, it has an atomic number of 19, meaning thatall potassium atoms and all potassium ions contain 19 protons.

Atomic Mass Number = (# of Protons) + (# of Neutrons)Example:

A particle with 6 protons and an atomic mass number of 14 has 8 neutrons.A particles name will sometimes include the atomic mass number of the particle.For example, chlorine-37 is a chlorine atom that has an atomic mass number of 37, meaning

that it has a total of 37 protons and neutrons in its nucleus.

Since the atomic number for chlorine is 17, any chlorine atom or ion always has 17 protons.Therefore, a chlorine-37 atom has 20 neutrons, because 37 minus 17 equals 20.

Ions = Charged ParticlesIons are formed when atoms gain or lose electrons.Positive ions (cations) are formed when a neutral atom loses electrons.

Negative ions (anions) are formed when a neutral atom gains electrons.Metallic atoms tend to lose electrons to form positive ions (also known as cations).

Nonmetallic atoms tend to gain electrons to form negative ions (a.k.a. anions).Charge = (# of Protons) - (# of Electrons)

Example:

A particle with 34 protons and 36 electrons has a charge of -2.

Isotopes

Main articles: Isotope and Stable isotope

Isotopes are atoms of the same element that have a different number of neutrons. Therefore,isotopes have the following characteristics:

Isotopes have the same atomic number (same number of protons), but a different atomic massnumber (a different number of neutrons). (Picture 1)

Isotopes behave the same chemically, because they are the same element. The only differenceis that one is heavier than the other, because of the additional neutrons.

For example, carbon-12 and carbon-14 are both isotopes of carbon. Carbon-12 has 6 neutrons;carbon-14 has 8 neutrons.The atomic weight of an element (as it appears in the periodic table) is the weighted average of

the atomic mass numbers of all of the isotopes for that element.For example, the atomic weight of carbon (as shown on the periodic table) is 12.011 a.m.u.

(atomic mass units). However, there is no one carbon atom that has a mass of 12.011 a.m.u.Carbon exists as four different isotopes: carbon-11, carbon-12, carbon-13, and carbon-14, whichhave approximate atomic mass numbers of 11, 12, 13, and 14 a.m.u., respectively.


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Pic-

ture

1.


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TOPIC 2: Biogenic s-elements, chemical properties, biological role and application

in medicine.

The s-block of the periodic table of elements consists of the first two groups: the alkali metalsand alkaline earth metals, plus hydrogen and helium.

These elements are distinguished by the property that in the atomic ground state, the highest-energy electron is in an s-orbital. Except in hydrogen and helium, these electrons are very easily

lost to form positive ions. The helium configuration chemically exceeds stablility and thus heliumhas no known stable compounds; thus it is generally grouped with the noble gases.

The other elements of the s-block are all extremely powerful reducing agents, so much that theynever occur naturally in the free state. The metallic forms of these elements can only be extracted

by electrolysis of a molten salt, since water is much more easily reduced to hydrogen than the ionsof these metals. Sir Humphry Davy, in 1807 and 1808, was the first to isolate all of these metalsexcept lithium, beryllium, rubidium and caesium. Beryllium was isolated independently by F.

All the s-block metals are dangerous fire hazards which require special extinguishers toextinguish. Except for beryllium and magnesium, storage must be under either argon or an inertliquid hydrocarbon. They react vigorously with water to liberate hydrogen, except for magnesium,which reacts slowly, and beryllium, which reacts only when amalgamated with mercury to destroy

the oxide film. Lithium has similar properties to magnesium due to the diagonal relationship withmagnesium in the periodic table.

The six elements belonging to group 1 of the periodic table, namely lithium (Li), sodium (Na),potassium (K), rubidium (Rb), cesium (Cs) and francium (Fr) are called alkali metals. Theyconstitute alkali metals because they readily dissolve in water to form hydroxides, which arestrongly alkaline in nature. They also form alkaline oxides. The francium element is radioactive.

Occurrence

Lithium occurs mainly as silicate minerals such as spodumene [LiAl(SiO3)2], lepidolite[(Li,Na,K)2Al2,(SiO3)3(F,OH)2] etc. It is the 35th most abundant element by weight.

Compounds of sodium and potassium have been known from ancient times. Sodium and

potassium are the seventh and eighth most abundant elements by weight in the Earth's crust andtogether make up over 4% of the Earth's crust by weight. NaCl and KCl also occur in large amountsin seawater. Rock salt (NaCl) is the major source of sodium. Potassium occurs mainly as deposits ofKCl (sylvite), which is a mixture of KCl and NaCl (sylvinite) and double salt KCl MgCl26H2O (carnallite).

Rubidium and cesium are obtained as a by-product of lithium processing. Francium beingradioactive does not occur appreciably in nature.

Nature of the compounds

The compounds of the alkali metals are ionic in nature. Alkali metals form cations readily bylosing the valence electrons (due to the low ionization energies and large atomic sizes). They go on

to form ionic bonds with the non-metals of the 'p' block.

Sodium carbonate

In anhydrous form, sodium carbonate is described by the chemical formula Na2CO3. Incrystalline state sodium carbonate exists as Na2CO3 10H2O (decahydrate) and is called washingsoda. It is a white crystalline solid, which loses water (effloresces) to dry air to yield monohydrate.

Na2CO3+ 10 H2O Na2CO3H2O + 9H2O

Sodium carbonate is used;1.As a filler in the detergent industry,

2.For washing purposes,3.For softening hard water,4.In the manufacture of many useful compounds such as, sodium hydroxide, borax (Na 2B4O7),

hypo (Na2S2O3) etc.5.In the paper and paint industry.


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Sodium bicarbonate is also known as baking soda. It's formula is NaHCO3. On heating, itdecomposes to give sodium carbonate and CO2is liberated.

2NaHCO3 Na2CO3+ H2O + CO2

Sodium bicarbonate is used for preparing baking powder and as a general treatment for acidityand mild indigestion i.e. as an effervescent drink (ENO salt).

The electronic configurations of alkali metals

Z Element Nom. of electrons

13

1119375587

HydrogenLithiumSodium

PotassiumRubidiumCaesiumFrancium

12, 12, 8, 12, 8, 8, 12, 8, 18, 8, 12, 8, 18, 18, 8, 12, 8, 18, 32, 18, 8, 1

Table 2.1Ionization energies

Alkali metals have the lowest ionization energy in each period. Within the group, as we godown, the ionization energies of alkali metals decrease due to their atomic size being the largest intheir respective periods. In large atoms the valence electrons are loosely held by the nucleus and areeasily lost, leading them to have low ionization energies and acquiring stable noble gasconfigurations. On moving down the group, the atomic size increases and the number of inner shellsalso increases, increasing the magnitude of screening effect and consequently, the ionization energydecreases down the group.

The second ionizations energies of alkali metals are very high. The removal of an electron fromalkali metals causes the formation of monovalent cations having very stable electronic

configurations (same as that of noble gases). Therefore, it becomes very difficult to remove thesecond electron from the stable noble gas configurations, giving very high second ionization energyvalues (IE2).(Table 2.2)

Table 2.2

Electro- positive or

metallic character

The electropos-itive character of an element is expressed in terms of the tendency of its atom to release electrons:

M M++ e-

All the alkali metals are strongly electropositive or metallic in character, since they have lowionization energies and their atoms readily lose the valence electron. As the ionization energiesdecrease down the family, the electron releasing tendency or electropositive character is expected toincrease down the family.

Oxidation states

All alkali metals have only one electron in their valence shell. They exhibit an oxidation state of+1 in their compounds and can lose the single valence electron readily to acquire the stable

configuration of a noble gas. Thus, they form monovalent ions, M+

(e.g., Li+

, Na+

, K+

, Rb+

, Cs+

).Since the second ionization energies are very high, they cannot form divalent ions. Thus, alkalimetals are univalent and form ionic compounds.

Characteristic flame coloration

Physical property Li Na K Rb CsIonization Energy

KJ mol-1III

5207298

4964562

4193051

4032633

3762230

Eo value (V) 3.03 - -2.71 -2.9

3 -2.93 - 2.92


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As the alkali metals have very low ionization energies, the energy from the flame of a bunsenburner is sufficient to excite the electrons of alkali metals to higher energy levels. The excited statebeing unstable, these electrons return to their original energy levels, emitting extra energy, whichgives characteristic flame colorations. The different colors of the alkali metals can be explained onthe basis of amount of energy absorbed for excitation of the valence electron.

Physical

property Li Na K Rb Cs

Flame colour crimson red Yellow pale violet violet bluishTable 2.3

Properties of Alkali Metals:

Table: Valence electron number of Alkali Metals

Hydrogen Lithium Sodium Potassium Rubidium Cesium Francium1 1 1 1 1 1 1

Table 2.4

Alkali metals are in the same group of the Periodic Table, and thus they also show a regulargradation in physical properties like boiling and melting points.

Thus Alkali metals are a group of metals that have almost same physical and chemical

properties:

Table: Properties of Alkali metals

Property Alkali Metals

Occurrence Combined state (in compounds)

Physical state Shiny, white, solid metals. Can be cut with a

knife.Valence eletrons One valence electron in each alkali metal.

Electrical Conduction Good conductors

Nature Electropositive; metallic nature increase down

the group.Reactivity Highly reactive

Reaction with water/acids Liberate Hydrogen on reacting with water and

acids; react vigorously and explosively.Types of compounds formed Electrovalent (Ionic) compounds

Atomic size Increases down the group.

Ionization energy Lowest ionization energy (because they lose

electrons).

Reducing/Oxidizing agents Strong reducing agents.

Table 2.5

Physical Characteristics of Alkali Metals- the alkali metals are very reactive elements and not found in free state in nature. Hence alkalimetals are normally stored in kerosene oil. Alkali metals are only isolated by electrolysis oftheir molten salts.The main physical characterestics of alkali metals are as:

- all alkali metal are soft metals with low density.- the colour of most alkali metals are silvery white, but Caesium is a golden colour metal.- alkali metal's boiling and melting point decreases while we move in alkali group from top to

bottom (Li have highest melting and boiling point).- alkali metals have low first ionization energy compared to all other group elements so due to

this all alkali metals shows its specific flame color like Sodium flame color is yellow.- all alkali metals are very good conductors of heat and electricity.

Chemical Characteristics of Alkali Metals


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All alkali metals are the most electropositive and least electronegative elements, they react withall types of nonmetals. Alkali metals are strong reducing agents. The value of standard electrode

potentials of all alkali metals are negative, which indicates its strong tendency to form cations like(Li+, Na+ , K+ , Rb+, Cs+ and Fr+) in solution. The main chemical properties of alkali metals aregiven below.

All alkali reacts with oxygen to form their coresponding oxide. Like when Li reacts with O2thenLi2O forms.

4Na + O2 2Na2O

Few alkali metals also form peroxides and superoxides.

2K2O + O2 2K2O2

Aklali metals reacts with halogen to form alkali halides (eg: NaCl, KBr, LiF, CsI etc.).

2Na + I2 2NaI

Alkali metals react with water very rapidally, When an Alkali metal reacts with water (H2O) thentheir corresponding alkali hydroxide form.

Li2O + H2O 2LiOH

Alkali metals are ionic compounds (except LiF, it is covalant compound because its cation size issmall).

Action with Hydrogen

Formation of hydrides

All alkali metals react with hydrogen to form hydrides that are ionic in nature (M+H-).

2M + H2 2M+H- (M = Li, Na, K, Rb, or Cs)

- reactivity of alkali metals with hydrogen increases from Li to Cs.

- ionic character of the hydrides increases from Li to Cs. The decrease in ionization energy downthe group permits easy availability of electrons to hydrogen, forming H- ion.

- stability of hydrides decreases from Li to Cs because the M-H bond becomes weak as the sizeof the alkali metal increases from Li to Cs. This causes the stability of hydrides to decrease.

- the hydrides behave as strong reducing agents and their reducing nature increases down thegroup.

Action of Halogens

Formation of halides

Alkali metals combine readily with halogens to form ionic halides M+X-. For example,

2M + X2 2MX (X = halogen)

2Na + Cl2 2NaCl

-reactivity of alkali metals with halogen increases down the group because of correspondingincrease in electropositive character (decrease in ionization energy).

-all metal halides are ionic crystals. Lithium iodide is slightly covalent as it has the smallestcation, which exerts maximum polarizing power and iodide ion being the largest anion can be

polarized to the largest extent.-all alkali metal halides are soluble in water except LiF, as it has a high lattice energy combined

with a small cation and anion, making it insoluble in water.

Properties of Alkali Metals and Halogens

Alkali metals:

alkali metals have less density compared to other metals.


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they have loosely bound valence electrons.they contain large radii in the periodsthey have low ionization energy and electronegativity.they are highly reactive in nature.

Action with Air

Formation of oxides

All the alkali metals on exposure to air or oxygen burn vigorously, forming oxides on the surfaceof the metals. Lithium forms monoxide (Li2O), sodium forms the peroxide (Na2O2) and the otherelements form superoxides.

(MO2: M = K, Rb, Cs).

4Li + O2 Li2O (monoxide)

2Na + O2 Na2O2 (peroxide)

M + O2 MO2 (0superoxide) (M = K, Rb, Cs)The alkali metal oxides are basic in nature because they dissolve in water to form alkali metal

hydroxides.

Na2O(s) + 2H2O(l) 2NaOH(aq)

Peroxides give hydrogen peroxide also

Na2O2(s) + 2H2O(l) 2NaOH(aq) + H2O2(aq)

Action Towards Water

Formation of hydroxides

Hydroxides and hydrogen gas result from the reaction of alkali metals with water.

2Na + 2H2O 2NaOH + H2(g)

2K + 2H2O 2KOH + H2

Hydration of ions

The alkali metal ions are highly hydrated. The smaller the size of the ion, the greater is thedegree of hydration. Thus, Li+ ion gets much more hydrated than Na+ ion which in turn is morehydrated than K+ ion and so on. The extent of hydration decreases down the group.

As a result of larger hydration of Li+ ion than Na+ ion, the effective size of Li+ ion is more thanthat of Na+ ion. Further the ionic radii in water (called hydrated ionic radii) decreases in the order:

Li+

> Na+

> K+

> Rb+

> Cs+ (

Table 2.6)

Ion Li+ Na+ K+ Rb+ Cs+

Ionic radius (pm) 76 102 138 152 167Hydrated radius (pm) 340 276 232 228 226Ionic mobility (ohm-1 cm2 mol-1) 33.5 43.5 64.5 67.5 68.0

Table 2.6

As a result, the hydrated Li+ ion being largest ionic size, has the lowest mobility in water. On theother hand, the hydrated Cs+ ion being smallest in size has the highest mobility in water.

Uses of Alkali Metals:

Lithium metal is used to make useful alloys, for example with lead to make white metalbearings for motor engines, with aluminium to make aircraft parts, and with magnesium to makearmour plates. It is used in thermonuclear reactions. Lithium is also used to make electrochemicalcells. Sodium is used to make a Na/Pb alloy needed to make PbEt4 and PbMe4. These organolead


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compounds were earlier used as anti-knock additives to petrol, but nowadays vehicles use lead-freepetrol. Liquid sodium metal is used as a coolant in fast breeder nuclear reactors. Potassium hasavital role in biological systems. Potassium chloride is used as a fertilizer. Potassium hydroxide isused in the manufacture of softsoap. It is also used as an excellent absorbent of carbondioxide.Caesium is used in devising photoelectric cells.

Uses of Group 1 Elements: Alkali Metals and Compounds: -1) Sodium chloride is used as common salt.2) Sodium is used in the preparation of Borax.3) Potassium bromide is used in photography.4) Potassium is used in the manufacture of NPK fertilizers.5) Potassium hydroxide and sodium hydroxide is used in the manufacture of detergents.6) Potassium chlorate is used in production of explosives.7) Rubidium and Francium are radioactive and are used in chemical researches.8) Sodium and Potassium are the vital nutrients in the human body. They help in Na- K pump.9) Sodium- bicarbonate is used in baking and is also known as baking powder10) Sodium compounds are also used as preservatives Ex:- MSG(monosodium glutamate).

Introduction to alkali earth metals properties:

The elements of 2nd A group of the periodic table are called the alkaline Earth metals. These arecalled alkaline earth metals because the oxides of these bases are alkali and they are found in theearth crust. Below are the following alkaline earth metals;

Be (4), Mg (12), Ca (20), Sr (38), Ba (56), Ra (88)

Electronic Configuration:

These elements have two electrons in the s -orbital of the valence shell. Their general electronicconfiguration may be represented as [noble gas] ns2. Like alkali metals, the compounds of theseelements are also predominantly ionic. (Table 2.7)

Table 2.7

Atomic and Ionic Radii:

The atomic and ionic radii of the alkaline earth metals are smaller than those of thecorresponding alkali metals in the same periods. This is due to the increased nuclear charge in theseelements. Within the group, the atomic and ionic radii increase with increase in atomic number.

Ionization Enthalpies:

The alkaline earth metals have low ionization enthalpies due to fairly large size of the atoms.Since the atomic size increases down the group, their ionization enthalpy decreases. The first

ionization enthalpies of the alkaline earth metals are higher than those of the corresponding Group 1metals. This is due to their small size as compared to the corresponding alkali metals. It isinteresting to note that the second ionization enthalpies of the alkaline earth metals are smaller thanthose of the corresponding alkali metals.

Physical Properties:

The alkaline earth metals in general are silvery white, lustrous and relatively soft but harder thanthe alkali metals. Beryllium and magnesium appear to be somewhat grayish. The melting and

boiling points of these metals are higher than the corresponding alkali metals due to their smaller

Element Symbol Electronic configuration:

Beryllium Be 1s22s2

Magnesium Mg 1s22s22p63s2

Calcium Ca 1s22s22p63s23p64s2

Strontium Sr 1s22s22p63s23p63d104s24p65s2

Barium Ba 1s22s22p63s23p63d104s24p64d105s25p66s2

Radium Ra [Rn]7s2


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sizes. The trend is, however, not systematic. Because of the low ionization enthalpies, they arestrongly electropositive in nature. The electropositive character increases down the group from Beto Ba. Calcium, strontium and barium impart characteristic brick red, crimson and apple greencolours respectively to the flame

Chemical Properties of Alkaline earth metals:The alkaline earth metals are less reactive than the alkali metals. The reactivity of these elements

increases on going down the group.1) Reactivity towards air and water:

Mg + 2H2O Mg(OH)2+ H2

2) Reactivity towards the halogens:

M + X2 MX2 (X = F, Cl, Br, l)

3) Reactivity towards hydrogen:All the elements except beryllium combine with hydrogen upon heating to form their hydrides,MH2. BeH2, however, can be prepared by the reaction of BeCl2with LiAlH4.

4) Reactivity towards acids:The alkaline earth metals readily react with acids liberating dihydrogen.

M + 2HCl MCl2+ H2

Uses of Bleacing Powder

It is used;1) as a disinfectant and germicide especially in the sterilization of drinking water.

2) for manufacture of chloroform.3) for making wool unshrinkable.4) as an oxidising agent in industry.5) mainly as bleaching agent for cotton, linen and wood pulp.However, delicate articles like straw, silk, ivory, etc., are not bleached by bleaching powder.Carbonic acid is a dibasic acid that gives rise to two series of salts, carbonates (normal salts) and

bicarbonates (acid salts), due to successive removal of the replaceable hydrogens from H2CO3.

H2CO3 NaHCO3+ H2O

NaHCO3+ NaOH Na2CO3+ H2O

Bicarbonates of calcium and magnesium are responsible for temporary hardness of water.

Preparation

By passing carbon dioxide through an alkali

With excess of carbon dioxide, the carbonate first formed changes to bicarbonate

With NaOH

2NaOH + CO2 Na2CO3+ H2Osodium carbonate

Na2CO3+ H2O + CO2 2NaHCO3sodium bicarbonate

With Ca(OH)2


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Ca(OH)2+ CO2 CaCO3+ H2O

CaCO3+ H2O + CO2 Ca(HCO3)2

By precipitation

Heavy metal carbonates are precipitated from their salt solutions with washing soda.

BaCl2+ Na2CO3 BaCO3+ 2NaCl

Bicarbonates of only alkali metals exist in solid state, even though carbonates of many metalsare known.

Bleaching powder is also called calcium chlorohypochlorite because it is considered as a mixedsalt of hydrochloric acid and hypochlorous acid. It is represented as:

PreparationIt is manufactured by the action of chlorine on dry slaked lime, Ca(OH) 2at 40oC.

Ca(OH)2+ Cl2 Ca(OCl)Cl + H2O

This is the Odling view about its formation. There is another view proposed by Cliffordaccording to which bleaching powder is a mixture of calcium hypochlorite and basic calciumchloride.

2Ca(OH)2+ 2Cl2 Ca(OCl)2+ CaCl2+ H2O

CaCl2+ Ca(OH)2+ H2O CaCl2Ca(OH)2H2O

2Cl2+ 3Ca(OH)2 Ca(OCl)2+ CaCl2Ca(OH)2H2O + H2O

The manufacture of bleaching powder is carried out in Backmann's plant as follows:It consists of a vertical cast-iron tower. The tower is provided with a hopper at the top, two inlets

near the base (one for chlorine and other for hot air) and an exit for waste gases near the top.The tower is fitted with eight shelves at different heights each equipped with rotating rakes.The slaked lime is introduced through the hopper and it comes in contact with chlorine, which

slowly moves upwards. Bleaching powder is collected in a barrel at the base. The chlorine used inthe manufacture of bleaching powder should be dilute and the temperature should be maintained

below 40oC.

Biological Importance of Sodium and Potassium

Sodium cation (Na+) is an important cation of extracellular fluids of animals including humanbeings which is known to mediate some enzymes in the animal body as well as in plants

Sodium ions are relatively harmless except when present in large quantities of amounts in whichcase they might cause hypertension in the environment. The saline water containing excessiveamounts of sodium chloride is harmful or injurious to plants and aquatic life because of the toxicityof Na+ ions.

The potassium ion is essential to all organisms with the possible exception of blue green algae. Itis a major cation of intracellular fluid of animal cells. It is moderately toxic to mammals but onlywhen injected intravenously. It acts as a catalytic enzyme.

Sodium pump/ Na+ - K+ pump


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The sodium (Na+) and potassium (K+) ions, taken together, constitute what is known as Na+ -K+ pump or simply as sodium pump. The sodium pump mechanism plays an important role in thesalinity tolerance in plants. Calculating sodium and potassium ions in plants is an important factorin salt tolerance. The function of this pump is genuinely a biological process which goes onoccurring in each and every cell of all animals. Although the discussion of the exact mechanism ofthe functioning of this pump is beyond the scope of the review yet it may be adequate to stay that ittransfers Na+ ions from the intracellular fluids to the extracellular fluids with the aid of proteins

known as carrier proteins.Simultaneously it transfers potassium K+ ions from the extracellular fluids to the intracellular

fluids. Since each operation of the pump (i.e. each cycle of this biological process) pumps out largernumber of sodium ions from the cell than the quantity of K+ ions that it pumps into the cell, theinterior of the cell acquires an excess negative (-ve) charge and the exterior of the cell acquires anexcess positive charge (+ve).

This results in the development of electrical potentialgradient across the cell membrane which is responsible for thetransmission of nerve signals in animals. The Na+ -K+ pumpalso accommodates the volume of the cell. Without the being

of the Na+-K+ pump in the cell, the cell would have swelled involume and finally brusted.

Picture 1.

The biological role of S elements


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Table 2.8

Element Location and role in the body Herbal drugs

Toxic

effect,

antidotes

H Element organogen

H2O23%antiseptic; a localhaemostatic;

HCl8,2-8,3%with reducedgastric acidity.

-

Na Extracellular cation. Buffer systems,

smosis, K, Na - pump

NaCl0,9%saline (isotonicsolution)a simple blood

substitution; for the preparation ofmedicinal substances;

NaCl4-10% hypertonicsolution;

NaHCO3baking soda, antacid;Na2SO4lenitive;

-

CaBone and dental tissue in the form of

compounds:a5(OH)(PO4)3orCaCO33Ca3(PO4)2H2O

CaCl2antiallergic, anti-inflammatory drug, increases

blood clotting.Cagluconateanti-inflammatory

effect;2CaSO42H2Oburnt plaster

casts;

-

Mg Intracellular ion; action against the

spasm

MgSO425%solution, a strongpurgative;

MgOmagnesia, the antacideffect;

MgCO3Mg(OH)23H2Owhitemagnesia, the antacid effect;

3MgO4SiO2H2Otalcumpowder, adsorbing agent for

powders;

-

Ba Retina Ba(SO4)2contrast agent in X-

ray

SolublesaltsBa2+

are toxic;antidotesNa2SO4,MgSO4


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TOPIC 3: Biogenic -elements, chemical properties, biological role,

application in medicine.

The elements belonging to groups 3 to 8 of the periodic table are the p-block elements.Elements of these groups have configuration ns2np16 (except He). The absence ofdorbitals in

second period and presence of d or f in heavier elements have significant effects on theirproperties. Apart fromthis the presence of all the three types of elements metals, metalloids and

nonmetals bring diversification in the chemistry of these groups.The p-block includes all the noble gases except helium, allthe nonmetals except hydrogen, all the

metalloids, and even a few metals, including Al, Sn, and Pb.Three of the p-block elements - O, Si,and Al - are the mostabundant elements in Earths crust. Six p-block elements - C, N, O, P, S, andCl - are among theelements making up the bulk of living matter.Five others - B, F, Si, Se, and I - arerequired in traceamounts by most plant and animal life.C and S can occur in the free state.

p - the elements are called chemical elements, which is filsl p-sublevel outer energy level. Theseare the elements III-A, IV-A, V-A, VI-A and VII-A groups of the periodic system ofD.I.Mendeleev. Electronic configuration, i.e. distribution of electrons in energy levels andsublevels, for atoms and ions of p - elements can be illustrated by P and Br:

P 1s22s22p63s23p3 P-3 1s22s22p63s23p 0

Br 1s22s22p63s23p63d104s24p5 Br-1 1s22s22p63s23p63d104s24p6.Some of these families from these group you can find here!

The nitrogen familyThe element nitrogen (N), phosphorus (P), arsenic (As),antimony (Sb) and bismuth (Bi)

constitute group 15 of theperiodic table.The first two elements of this group, nitrogenandphosphorus are non-metals, the next two, arsenic antimonyare metalloids whereas the lastelement, bismuth, is a metal.All the elements of this group have five electrons in theirvalenceshells.Their general valence shell electronic configuration isns2np3.

Physical properties

Melting points and boiling points.These increase on going down the group. However, the melting point of bismuth is usually low.

This may presumably be due to more tendency of Bi toform three rather than five covalent bonds(inert paireffect). Thus, there are weak forces of attraction between theiratoms in solid state,accounting for its low m.pt.

Atomic radii.

Atomic radii, as expected, increases with increase inatomic number.

Ionization energy and electropositive character.

The first ionisation energy decreases regularly down thegroup due to increase in the size ofatoms andscreening effect of intervening electrons.The ionisation energies of this group are muchhigherthan those of corresponding elements of group 14.This is due to the increase in nuclearcharge and extrastable configurations of the elements of this group.The extra stability ofconfigurations is attributed to the exactly half-filled p-orbitals in their valence shells.Due to highionisation energies, the elements of this group are less electropositive than elements of group14.Electropositive (or metallic) character increases on going down the group due to decrease inionisation energies.The fact is reflected by the change in metallic characterfrom nitrogen to

bismuth.

Oxidation States.

All the elements of this group have five electrons in their valence shell.These elements requirethree more electrons to acquire noble gas configuration. However, though gaining of three electrons


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requires alarge amount of energy yet it takes place only withnitrogen as it is the smallest and themost electronegative element of this group.N3ion exists in the compounds such as Mg3N2, Ca3N2,etc.The other elements of the group form covalent compounds in which their formal oxidation stateis 3.For example, calcium phosphide (Ca3P2), sodium arsenide (Na3As), zinc antimonide (Zn3Sb2)and magnesium bismuthide (Mg3Bi2).In addition to 3 oxidation state, the elements of this groupalso exhibit + 3 and + 5 oxidation states.

Chemical propertiesHydrides

The elements of group 15 form hydrides having the general formula MH3. All these are covalentin nature.

1) The thermal stability of the hydrides decreases in the order:NH3> PH3> AsH3> SbH3> BiH3

It is because the strength of MH bond decreases down the group due to increase in the size ofcentral atom.

2) These hydrides act as reducing agents.The least stable hydride (BiH3) acts as the strongestreducing agent whereas the most stable hydride (NH3) acts as a weakest reducing agent.

3) The boiling points of these hydrides vary as follows:

Hydride NH3 PH3 AsH3 SbH3 BiH3Boiling point (K) 238.5 185.5 210 254 290

NH3 molecules are associated by intermolecular Hbonds.As a result its boiling point isexceptionally high.The interparticle forces in PH3are van der Waal forces due to which its boiling

point is low.In moving from PH3 to BiH3 boiling points increase.This is due to the increase in themagnitude of van der Waal forces due to increase in molecular size.

4) Due to presence of a lone pair of electrons on the central atom in these hydrides, these act asgood Lewis bases.The basic character decreases down the group.It is because of small size ofnitrogen. The lone pair of electrons are concentrated on a small region so density per unit volume ismore.

Halides

All the elements of group 15 form two series of halides, i.e., trihalides and pentahalides of thetype MX3 and MX5.These trihalides are mainly covalent with the exception of BiF3 which isionic.The ionic character of trihalides increases in going down the group.These trihalides except

NX3can be easily hydrolysed by water.The inability of trihalides of N to hydrolyse is attributed tothe non-availability of vacant d-orbitals in nitrogen.

The trihalides of P, As, Sb (especially fluorides and chlorides) act as Lewis acids and combinewith Lewis bases). Except nitrogen other elements of this group form pentahalides. Nitrogen cannotform pentahalides because it cannot expand its covalency due to non-availability of d-orbitals in thevalence shell.All the pentahalides act as Lewis acids.It is because the central atom can easily acceptthe halide ions (due to presence of vacant d-orbital) and can extend their co-ordination number.

Oxides

The elements of this group combine with oxygen directly or indirectly to form a large number ofdifferent types of oxides. All those elements form 2 types of oxides E 2O3and E2O5.All the oxides ofnitrogen (except NO and N2O) and phosphorus are strongly acidic.Oxides of arsenic are weaklyacidic; oxides of antimony are amphoteric and those of bismuth are weakly basic.As far as thestability of the oxides is concerned, it is found that oxides having elements in the higher oxidationstate become less stable as we move down the group.This is because of the inert pair effect.

NitrogenNitrogen is found in greater abundance in the atmosphere than anywhere else.There are only twoimportant mineral sources of nitrogen: KNO3 and NaNO3.Nitrogen compounds occur in all livingmatter.Nitrogen molecule, N2, has a very strong N N triple bond. Consequently it is quite


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unreactive and is used as an inert blanketing atmosphere in industrial operations.Liquid nitrogen isused in low-temperature applications.The only important commercial method of producing nitrogenis the fractional distillation of liquid air.

Preparation of dinitrogen

In laboratory, dinitrogen is prepared by heating aqueous solution of ammonium chloride withsodium nitrite.

NH4Cl(aq) + NaNO2 N2(g) +2H2O(l) + NaCl(aq)

Small amount of NO and HNO3 are also formed in this reaction; these impurities can beremoved by passing the gas through aqueous sulphuric acid containing some potassiumdichromate.

Properties of dinitroge

Some of the physical properties of dinitrogen are:1) It is colourless, tasteless and odourless gas:2) It is non-toxic.3) It has very low solubility in water (23.2 xm3per litre of water at 273 K and 1 atm.)4) Its m.p. and b.p. are 63.2 K and 77 K respectively.

Chemical properties

Dinitrogen has a very little chemical reactivity at ordinary temperatures.It is neither combustiblenor supporter of combustion.It may be noted that nitrogen is neither strongly electronegative to actas good oxidising agent nor strongly electropositive to act as good reducing agent.

Reaction with highly electropositive metals like lithium, calcium, magnesium, etc.

6Li + N2 2Li3N

3Ca + N2 Ca3N2

Reaction with non-metals like dihydrogen and dioxygen

N2+ 3H2 2NH3Ammonia

N2+ O2 2NONitricoxide

Use of dinitrogen

The main use of dinitrogen is in the manufacture ofammonia and other industrial nitrogen

chemicals like calcium cyanamide.It also finds use where the presence of an inert gas is required.Liquid nitrogen is used as refrigerant to preserve biological materials, in freezing food articles

and in cryosurgery. It is also used in gas filled mercury thermometers employed for measuring hightemperatures.

Ammonia (NH3)

Preparation of ammonia

By heating ammonium salts.

(NH4)SO4 2NH3+ H2SO4Amm.sulphate

NH4NaHPO4 NH3+ NaPO2+ H2OAmm. sodium hydrogenphosphate Sod. metaphosphite


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When the ammonium salts of non-volatile acids such as ammonium sulphate, ammoniumsodium hydrogen phosphate or ammonium phosphate are heated, ammonia is produced.

By heating ammonium salts with a strong bases

NH4Cl + KOH KCl +H2O +NH3

(NH4)2SO4+ 2NaOH Na2SO4+ 2H2O + 2NH3

Large scale preparationAmmonia is generally manufactured by Habers process which involves the direct combination

of nitrogen and hydrogen.

N2(g) + 3H(g) 2NH3(g); =93.6kJ/mol .

Chemical properties

Basic NatureAqueous solution of ammonia is basic in nature and turns red litmus blue. Its basic character is

due to the formation of OH ions in aqueous solution.In gaseous state also, it reacts with acids to form salts

NH3(g) + HCl(g) NH4Cl(s)

2NH3(g) + H2SO4(l) (NH4)2SO4

Complex FormationAmmonia acts as a Lewis base due to the presence of a lone pair of electrons on nitrogen and

hence, forms a number of complexes with metal ions.Due to the tendency of complex formation thewhite precipitate of silver chloride dissolves in aqueous solution of ammonium hydroxide.

Precipitation of heavy metal ions from the aqueous solutions of their saltsHeavy metal ions like Fe3+, Al3+, Cr3+, etc., are precipitated from their aqueous salt solutions.

FeCl3+3NH4OH Fe(OH)3+ 3NH4Cl

AlCl3+ 3NH4OH Al(OH)3+ 3NH4Cl

Use of ammonia

Ammonia is used mostly to produce various nitrogeneous fertilizers (ammonium nitrate, urea,ammonium phosphate and ammonium sulphate). Liquid ammonia is also used as a refrigerant.

Nitrogen forms a number of oxides in different oxidation states

Details are given in the table 3.1

Name Formula

Oxidation

state of

nitrogen

Common methods of

preparation

Physical

appearance and

chemical nature

Dinitrogen oxide[Nitrogen(I) oxide]

N2O +1 NH4NO3N2O + 2H2O

colourless gas,neutral

Nitrogen monoxide[Nitrogen(II) oxide]

NO +2

2NaNO2+ 2FeSO4+

3H2SO4Fe2(SO4)3+ 2NaHSO4+

2H2O + 2NO

colourless gas,neutral

Dinitrogen trioxide[Nitrogen(IV) oxide]

N2O3 +32NO + N2O4

2N2O3blue solid, acidic


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Nitrogen dioxide[Nitrogen(IV) oxide]

NO2 +4

2Pb(NO3)2

4NO2+2PbO2

brown gas, acidic

Dinitrogen tetroxide[Nitrogen(IV) oxide]

N2O4 +4 2NO2N2O4

colourlesssolid/liquid,

acidic

Dinitrogen pentoxide[Nitrogen(V) oxide]

N2O5 +54HNO3+ P4O10

4HPO3+2N2O5

colourless solid,acidic

Table 3.1

Nitric acid (HNO3)

Manufacture (Ostwalds process).This method was established by Wilhelm Ostwald (1902). In this method, HNO3 is

manufactured by oxidation of ammonia.It takes place in following steps.

1) Catalytic oxidation of ammoniaA gaseous mixture containing pure and dry ammonia and dust free dry air in the ratio of 1:10 by

volume is passed through a steel shell or aluminium converter having platinum gauze heatedelectrically to about 1100 K.

Ammonia is oxidized to nitric oxide and the reaction is exothermic.

4NH3+ 5O2 4NO +6H2O

2) Oxidation of nitric oxideThe nitric oxide obtained above is cooled and passed through a spacious chamber called

oxidising chamber. The nitric oxide is oxidized to nitrogen dioxide.

2NO + O2 2NO2

3) Absorption towerThe gases are then introduced in absorption tower which is packed with acid proof stones.Water

sprayed from the top dissolves nitrogen dioxide in the presence of oxygen to give about 60% nitricacid.

3NO2+ H2O 2HNO3+ NO

The acid flows down and is withdrawn from the tower with the help of a tap. Furtherconcentration of nitric acid can be done to 98% acid by dehydration with concentrated sulphuricacid.

Laboratory preparation

In the laboratory, nitric acid can be prepared by heating NaNO3 or KNO3 with concentratedsulphuric acid in a glass retort and condensing the vapours of HNO3.

NaNO3+ H2SO4 NaHSO4+ HNO3

Physical properties of nitric acid

When pure, it is a colourless liquid.The impure acid is generally yellow due to the presence ofnitrogen dioxide as impurity.Nitric acid containing dissolved nitrogen dioxide is known asfumingnitric acid.It has corrosive action on skin and produces painful blisters.Pure acid has a specificgravity of 1.54. It boils at 359 K and freezes to a white solid (m.p. 231 K).

Chemical properties

Acidic nature.


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Nitric acid is one of the strongest acids. In aqueous solutions it gets almost completely ionised toform H3O+ ions and NO3ions.

HNO3(aq) + H2O(l) H3O+(aq) + NO3-(aq)

Therefore, nitric acid gives the usual reactions of acids.It turns blue litmus acid.It reacts withbases to form salts and water. For example:

NaOH(aq) + HNO3(aq) NaNO3(aq) + H2O(l)

It decomposes carbonates to give carbon dioxide.

Action on Metals

When nitric acid reacts with metals, different products are formed depending upon theconcentration of the acid and activity of the metal.

Dilute nitric acid

Magnesium and manganese are the only metals which liberate hydrogen from very dilute nitricacid (2%).

Mg(s) + 2HNO2(aq) Mg(NO3)2(aq) + H2(g)

Mn(s) + 2HNO2(aq) Mn(NO3)2(aq) + H2(g)

More active metals such as zinc, tin and iron, react with cold dilute nitric acid to formammonium nitrate.

4Zn(s) + 10HNO3(aq) 4Zn(NO3)2(aq) + NH4NO3(aq) + 3H2O(l)

4Sn(s) + 10HNO3(aq) 4Sn(NO3)2(aq) + NH4NO3(aq) + 3H2O(l)

However, with hot dilute nitric acid the ammonium nitrate, so formed gets decomposed resultingin the production of N2O.

4Zn(s) + 10HNO3(aq) 4Zn(NO3)2(aq) + N2O(g) + 5H2O(l)

Less active metals like lead, silver, mercury and copper, react with dilute nitric acid to liberatenitric oxide NO as:

3Pb(s) + 8HNO3(aq) 3Pb(NO3)2(aq) + 2NO(g) + 4H2O(l)

3Cu(s) + 8HNO3(aq) 3Cu(NO3)2(aq) + 2NO(g) + 4H2O(l)

Concentrated nitric acid

Metals such as lead, silver, copper and zinc, form nitrogen dioxide as:

P4+ 20HNO3 4H3PO4+ 20NO2+ 4H2O

Pb(s) + 4HNO3(aq) Pb(NO3)2(aq) + 2NO2(g) + 2H2O(l)

Metals, such as iron, aluminium and nickel are rendered passive by concentrated nitric acid.The passivity of these metals is attributed to the formation of thin protective layer of metal oxide

on the surface of metal which cuts off the further action. Noble metals like gold and platinum donot react with nitric acid.However they dissolve in aqua regia (mixture of one part of conc. HNO3

and 3 parts of conc. HCl) forming respective chlorides.

Phosphorus


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Phosphorus is the eleventh most abundant element in Earths crust occurs exclusively innature as phosphate: PO43.

Elemental forms:(I) White phosphorus, P4, can be cut with a knife, melts at 44.1 oC, is a non-conductor of

electricity, and ignites spontaneously in air (it is stored under water).(II) Red phosphorus can be obtained by heating white phosphorus to about 300 oC in the absence

of air. This allotrope of phosphorus forms long chains of phosphorus atoms joined together.

Phosphine

It is prepared by the following methods.

From phosphides.

By the reaction of phosphides with water.

Ca3P2+ 6H2O 2PH3+ 3Ca(OH)2Calcium phosphide Phosphine

Na3P + 3H2O PH3+ 3NaOH

Sodium phosphide

Laboratory method.

Phosphine is prepared in the laboratory by heating white phosphorus with concentrated sodiumhydroxide solution in an inert atmosphere of carbon dioxide or oil gas in a round-bottom flask.

P4+ 3NaOH + 3H2O 3NaH2PO2+ PH3Sod. hydrophosphite

Phosphine evolved is spontaneously inflammable due to the presence of phosphorus dihydride(P2H4, the liquid hydride) as impurity.

Properties

It is a colourless, poisonous gas with rotten fish smell. It is slightly soluble in water and itssolution in water decomposes in presence of light giving red phosphorus and H2. On being absorbedin copper sulphate or mercuric chloride solution, corresponding phosphides are obtained.

3CuSO4+ 2PH3 Cu3P2+ 3H2SO4

3HgCl2+ 2PH3 Hg3P2+ 6HCl

Like NH3it is weakly basic and gives phosphonium compounds with acids.

PH3+ HI PH4I

Phosphorus forms a number of oxo acids

Details are given in the table 3.2

Name

Hypophosphorous

acid (phosphinic)

Orthophosphorous acid

(phosphonic)

Ortho phosphoricacid

Pyro phosphoricacid

Formula H3PO2 H3PO3 H3PO4 H4P2O7Oxidation state of P +1 +3 +5 +5

Preparation of

acid

White P4+ alkali(Power reducing

agent)

P2O3+H2O(weak reducing

agent)

P4O10+ H2O(No reducing

properties)

Heat phosphoricacid

(No reducingproperties)Table 3.2

The Oxygen Family


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Group 16 of periodic table contains five elements namely, oxygen (O), sulphur (S), selenium(Se), tellurium (Te) and polonium (Po). Group 16 of periodic table contains five elements namely,oxygen (O), sulphur (S), selenium (Se), tellurium (Te) and polonium (Po).

These are collectively known as chalcogens or ore forming elements because many metal oresoccur as oxides and sulphides.

Oxygen, the most abundant elements, is an important constituent of atmosphere and ocean. Itconstitutes 46.6% of earths crust.

Oxygen

Oxygen is one of the most active non-metals and one of the most important. The chief reactionsof elemental, atmospheric oxygen are oxidation processes. Uses of oxygen include: manufacture ofiron, other metals, welding, manufacture of chemicals, water treatment, oxidizer, and respirationtherapy. Ozone, O3, is a powerful oxidizing agent, especially in acidic solution. It is also found inthe upper atmosphere.

Sulphur

Sulfur forms many compounds similar to those of oxygen.However They are differ in importantway: hydrogen bonding in O compounds, but not in S compounds and S can employ an expanded

valence shell, but O cannot.Elemental sulfur exists as several molecular species: solid - S8, vapor - S2. Elemental sulphur is

mined using the Frasch process. A small amount of sulphur is used directly in vulcanizing rubberand as a pesticide.


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Physical properties

Density

As expected, the density increases down the group with the increase in atomic number.

Melting and Boiling Points

The melting and boiling points show a regular increase down the group.It is because as we movedown the group the molecular size increases. As a result, the magnitude of van der Waal forces

increase and hence the melting and boiling points also increase.

Ionization energy

These elements possess large values of ionization energies which decrease gradually from O toPo. Oxygen has very high ionization energy (1314 kJ mol1). The decrease in ionization energyfrom oxygen to polonium is due to increase in size of the atoms and increase in screening effect ofthe electrons belonging to lower shells.

Metallic and non-metallic character

Metallic character depends on ionization energy. Lesser the ionization energy greater will be themetallic character. As we move down the group, the ionization energy decreases and therefore the

metallic character increase. Thus, the first four elements namely oxygen, sulphur, selenium andtellurium are non-metals. The non-metallic character is maximum in O and decreases down thegroup.Polonium, the last member of the group, is metallic in nature.

Oxidation states

All the elements have ns2np4electronic configuration of outermost shell in their atoms. Thus,these tend to acquire the noble gas configuration by sharing or gaining electrons. Oxygen, the firstmember of the family exhibits the oxidation state of 2 in its compounds owing to its highelectronegativity.In addition to the oxidation state of 2, oxygen also exhibits an oxidation state of1 in H2O2, zero in O2and + 2 in OF2.

Electronegativity

Oxygen is strongly electronegative in character. The value of electronegativity decreases withincrease in atomic number down the group. The value of electronegativity decreases with increasein atomic number down the group.

Chemical properties

Hydride

These elements form volatile hydrides of H2E type such as H2O, H2S, H2Se, H2Te and H2Po. Thevolatile nature of hydrides first increases from H2O to H2S and then decreases. The low volatilecharacter of water is due to the association of water molecules through hydrogen bonding. The

decrease in volatile nature from H2S onwards is due to increased interactions among molecules dueto increase in molecular size. These hydrides are weak lyacidic in nature. The acidic natureincreased from H2O to H2Te. H2S is a weak diprotic acid.

The increase in strength as acid on moving down the group is due to increase in size of theatoms. The distance between central atom and hydrogen increases as we move from O to Te whichfavours the release of hydrogen as proton.This order of acidity, however, cannot be explained interms electronegativity of these elements. All the hydrides, except water, are reducing agents. Thereducing character increase on moving from H2S to H2Te. This is presumably due to the increase inthe size of the central atom due to which the strength of MH bond and also the stability decreases down the group.

HalideThe elements of group 16 form a number of halides. The more important halides being hexa,

tetra- and dihalides. The stability of halides is in the following order F> Cl< Br > I. Thehexaflurorides of S, Se and Te i.e., SF6, SeF6, TeF6 have the central atom in sp3d2hybridised


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state.These fluorides have octahedral structure.It is the small size of fluorine which brings out themaximum covalency. Hexahalides with Cl, Br and I are not formed. Due to the bigger size of thesehalogen atoms a coordination number of six is not achieved. Amongst the tetrafluorides of S, Seand Te, SF4is gaseous SeF4is liquid and TeF4is solid. In tetrahalides sp3d hybridisation occurs andthus these have got a trigonal bipyramidal structure with one of the equatorial position occupied bya lone pair of election.

OxidesThe elements of this group form a number of oxides.All elements (except Se) form monoxide.

All elements form dioxide with formula MO2.SO2is a gas, SeO2is a volatile solid while TeO2andPoO2are non-volatile crystalline solids. The acidic character of the dioxides decreases as we movedown the group.All the group 16 elements form the trioxides, MO3. The best known trioxide is SO3.

Oxygen

Preparation of Dioxygen

From oxygen rich salts

Certain oxygen rich salts such as chlorates, permanganates, peroxides and nitrates onheatingdecompose to give dioxygen gas.

2KClO3(s) 2KCl(s) + 3O2(g)Potassium chlorate

3MnO2 Mn3O4+ O2Manganese dioxide

Manufacture of Dioxygen

From air

The most economical method for commercial preparation of dioxygen involves liquefaction ofair followed by fractional distillation of the liquid air thus obtained.

From water

Dioxygen can also be prepared by the electrolysis of water containing a small amount of amineral acid or an alkali.

2H2O(l) 2H2(g) + O2(g)

Properties of Dioxygen

1) Dioxygen is colourless, tasteless and odourless gas.2) It is slightly soluble in water.3) Its solubility being approximately 30 cm3per liter of water at 298 K.

4) Its slight solubility in water is responsible for sustaining the life of aquatic animals and alsofor the degradation oforganic wastes in water bodies.

5) It is appreciably soluble in alkaline pyrogallol solution.6) Liquid oxygen exhibits, paramagnetism, i.e., it is slightly attracted by magnet.

Chemical reactions

Action with litmus.

It is neutral to litmus.

Reaction with metals.

Active metals like Na, Ca, etc. react at room temperature forming the respective oxides.

4Na(s) + O2(g) 2Na2O2(s)Sodiumperoxide


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2Ca(s) + O2(g) 2CaO(s)

Metals like Mg burn in dioxygen to form magnesium oxide.

2Mg(s) + O2(g) 2MgO(s)

Metals like iron and aluminium combine with dioxygen on heating to form their respectiveoxides.

4Fe(s) + 3O2(g) 2Fe2O3(s)

Reaction with non-metals.

Phosphorous

P4(s) + 5O2 P4O10(s)

Dinitrogen

N2(g) + O2(g) 2NO(g)

Reaction with compounds.

Hydrogen chlorid

4HCl(g) + O2(g) 2H2O(l) + 2Cl2(g)

Sulphur dioxide

2SO2(g) + O2(g) 2SO3(g)

Sulphur dioxide (SO2)

Sulphur dioxide is formed by burning sulphur in air or roasting metal sulphides in the presenceof air.

S8+ 8O2 8SO2

4FeS2 + 11O2 2Fe2O3+ 8SO2

Sulphuric acid (H2SO4)

It is the most important oxyacid of sulphur and is known from ancient times.It is calledoil ofvitriolbecause in early days it has been prepared from ferrous sulphate crystals (green vitriol) andhas an oily appearance. Because of its large applications in industries it is also known as thekingof chemicals.

Manufacture of sulphuric acidSulphuric acid is manufactured by contact process. The process involves the following steps:

Production of sulphur dioxide.

It is carried out by burning powdered sulphur or roasting of sulphur rich ores.

S8+ 8O2 8SO2

4FeS2+ 11O2 2FeO3+ 8SO2

Conversion of SO3into H2SO4.SO3is absorbed in conc. H2SO4to get oleum.

SO3+ H2SO4 H2S2O7


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Oleum may then be diluted with calculated quantity of water to get H2SO4of requiredconcentration.

H2S2O7+ H2O 2H2SO4

The HalogensHalogens are all non-metals with ns2np5 valence shell electron configuration. Halogen elements

exist as diatomic molecules, X2. Fluorine is the most reactive, iodine is the least. The oxidizing

power decreases from F2to I2. Halogens occur naturally only as the halides (X) and are convertedto halogens usually by electrolysis.

Physical properties

Melting and Boiling Points.

Melting and boiling points increase with increase in atomic number. The heat of fusion as wellas heat of vapourisation also increase with the increase in atomic number. This indicates that thestrength of intermolecular forces of attraction between the molecules increases with the increase inatomic number.

Atomic and Ionic Radii

Atomic and ionic radii are small and increase regularly down the group form fluorine to iodinebecause new electronic shells are added.

Ionization Energies

Ionization energies of all the halogens are very high. Therefore, they have a less tendency to loseelectron. However, this tendency increases down the group because the distance of valence shellfrom nucleus increases. Iodine is capable of forming stable compounds in which it exists as I - ion.

Colour

Halogens are coloured.The colour of different halogens are as follows:

Details are given in the table 3.3Halogen Fluorine Chlorine Bromine Iodine

Colour Pale Yellow Greenish Yellow Reddish brown Dark violet

Table 3.3

Chemical propertiesAll the halogens are very reactive but amongst them fluorine is the most reactive.As we move

down the group, reactivity decreases. This is due to the decrease in electronegativity and increase inbond dissociation energy. In general, a halogen of low atomic number oxidizes halide ions of higheratomic number i.e., F2displaces Cl2, Br2and I2from their salts, Cl2displaces Br2 and I2where asBr2 displaces iodine from its salts, e.g.,

F2+ X- 2F- + X2 (X = Cl, Br, I)

Cl2+ 2X- 2Cl- + X2 (X = Br, I)

Oxidising power

The halogens act as strong oxidising agents. It is because of their high electron affinities thatthey have great tendency to take up electrons. The oxidising power of the halogens decreasesongoing down the group from fluorine to astatine.

Hydride

Halogens combine with hydrogen to form volatile halides of the formula HX. Except HF, allother hydrides HCl, HBr and HI are gases. HF is a liquid because of intermolecular hydrogen

bonding.

HF.HF.HF.


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All the hydrogen halides act as acids in their aqueous solutions. The acidic strength varies in theorder.

HF < HCl < HBr < HI

Halides

The halogens, in general, react with almost all the metals and non-metals except, He, Ne and Arto form a wide variety of binary halides. With metals like Na, K, Mg, etc. having low ionisation

energies, the halogens react to form ionic halides which have got high boiling and melting points.As the electronegativity of the halogens decreases down the group the ionic character decreases inthe order.

MF > MCl > MBr > MI

Oxides

Halogens do not combine with oxygen directly but their oxides with oxygen can be preparedindirectly. For example, oxygen difluoride, OF2is prepared by the action of fluorine on 2% sodiumhydroxide solution. The compounds of oxygen with fluorine are called as fluorides because fluorineis more electronegative than oxygen. Most of these compounds are endothermic and unstable andare likely to explode resulting in the formation of more stable products.

The biological role of Pelements

Table3.4

Element Location and role

in the body Herbal drugs

Toxic effect,

antidotes

B Carbohydrate-phosphorus

metabolism

H3BO3disinfectant properties (eyeand ear drops);

Na2B4O7(bur)antiseptic-

Alblood, nerve cells in the brain;

involved in the construction of theepithelial and connective tissue

Al(OH)3absorbent and antacidproperties;

almagel - water suspension;

Al2O32SiO22H2Okaolin,adsorbing action;Al2(SO4)3 hemostatic,

antimicrobial action and for waterpurification.

KAl(SO4)212H2O (lum)hemostatic, antimicrobial action;

-

C rganogen, 21,15%

(arbol, activated charcoal)flatulence adsorb gases, toxic

substances;O2stimulatory effect on

respiratory centers, inhalations,baths;

NaHCO3baking soda, antacid.

Coal dustanthracosis;CO2carbon

monoxide;antidote -xygen

Si Lens of the eye, hair; gives

strength, elastic fabric.Silicon carbide and oxide used in

dentistry.SiO2dust causes

silicosis.

Pb The biological role has not been

studied

(CH3COO)(OH)Pb lead water,anti-inflammatory, antimicrobial

action.

Pb2+ toxic, bindsSH-groups of

proteins, enzymes;

N rganogen; 3,1%; proteins,

nucleic acids.

NH4OH9,5-10,5% solution,irritating effect on the CNS;

NH4Cl diuretic;

NaNO2 vasodilator;N2O inhaled anesthetics.

-

P

rganogen; 0,95% nucleic acids,ATP,bone and dental tissue in the

form ofcompounds:a5(OH)(PO4)3or

CaCO33Ca3(PO4)2H2O

Calcium glycerophosphatea meansof fortifying;

ATPthe energy product.

honky poison;antidote0,5%solutionuSO4

As Brain tissue, muscle, involved in

the synthesis of hemoglobin.

Asrganic compoundsfor thetreatment of sexually transmitted

diseases;

As2O3necrotising tissue (used indentistry)

As2O3 whitearsenic, a powerfulpoison antidotes

Na2S, MgS,Na2S2O3

O rganogen; 62,4%; O2 + CO2stimulates the

respiratory center; -

Srganogen; 0,16%; proteins,

amino acids - cysteine,methionine;

S (leaned)antimicrobial action;

SO2disinfectant ;Na2SO4weakpurgative;

S rganic compounds - sulfa drugsantimicrobialaction ;

H3C-SO-CH3dimixed; well

penetrates through biologicalmembranes, anti-inflammatory

effect.

SO2 irritating tomucous membranes

of the respiratorytract and eyes.

F Bone and dental tissue Ca5(PO4)3F NaF, KF,sedativesExcess fluoride

causesfluorosisorspeckled enamel

Cl Gastric juice, extracellular anion HCl8,2-8,3%at low acidity of Cl2gas, irritating


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TOPIC 4: Biogenic d- elements, chemical properties, biological role,

application in medicine.

The elements which have partially filled d-orbitals either ground state or in one or more of theirions, are called d-block elements or outer transition elements. Their properties are intermediate

between s-block elements and p-block elements. They are more electropositive than p-blockelements but less electropositive than s-block elements. They form ionic compounds in their lower

oxidation state and covalent compounds in higher oxidation states. They are all metals and haveelectronic configuration ns2, (n-1)d1 to 10

General characteristics of d-block elements

Electronic configuration of first series of d-block elements. (Table 4.1)

For first series (Sc z = 21) to (Zn z = 30)

Sc (z = 21) 1s2 2s2 2p6 3s2 3p6 4s2 3d1

Ti (z = 22) 1s2 2s2 2p6 3s2 3p6 4s2 3d2

V (z = 23) 1s2 2s2 2p6 3s2 3p6 4s2 3d3

Cr (z = 24) 1s2 2s2 2p6 3s2 3p6 4s1 3d5

Mn (z = 25) 1s2 2s2 2p6 3s2 3p6 4s2 3d5

Fe (z = 26) 1s2 2s2 2p6 3s2 3p6 4s2 3d6

Co (z = 27) 1s2 2s2 2p6 3s2 3p6 4s2 3d7

Ni (z = 28) 1s2 2s2 2p6 3s2 3p6 4s2 3d8

Cu (z = 29) 1s2 2s2 2p6 3s2 3p6 4s1 3d1

0

Zn (z = 30) 1s2 2s2 2p6 3s2 3p6 4s2 3d1

0

Table 4.1

Metallic character

All transition elements are metals.

Atomic size

The atomic size of d-block elements decreases from Sc to Zn due to increase in nuclear charge.Their atomic sizes are smaller than that of s-block elements.

Ionization potential

The I.P. values of transition elements are intermediate between those of s-block and p-blockelements. This shows that they are more electropositive than p-block elements and less reactivethan s-block elements. Due to intermediate values I.P. they form ionic compounds as well ascovalent compounds.

Melting and boiling point

Melting and boiling point of these elements are very high (except Zn). These higher values aredue to small atomic radii of transition elements which provides greater inter atomic forces ofattraction. They are very hard .

Conductivity

Transition elements are good conductor of heat and electricity.

Colour

Except zinc all the transition metals complex ions are colourless due to presence of unpaired

electrons. The colour of ions can be explained on the basis of "Crystal field theory". According tothis theory ,the bonding between ligands and a metal ion is electrostatic. The ligands surroundingthe metal ion and create an electrostatic field around its d-orbitals. This field split '5' degenerated d-orbitals in to two sets of different energies.


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a) A high energy pair (eg) of dx2-dy2 and dz2.b) A low energy trio (t2g) of dxy, dyz, dzx.In many cases difference of energy between two sets 'eg' and 't2g' is equivalent to a wavelength

in the visible region. Thus absorbing visible light, an electron may be able to move from lowerenergy set t2g to higher energy set eg. In doing so, some of the component wavelength of whitelight is removed, so the remaining component wavelength of light reflected or transmitted showsthe colour.

For example: Cu+2

(blue), V+3

(green), Co+2

(pink), Fe+2

(green), Fe+3

(yellow), Ti+3

(purple).At different oxidation states, same element produce different colours.

General Features

The d-block elements is the collective name for groups 3 to 12 in the periodic table.They are so called because a d-subshell is being filled. Here only the first row, from scandium to

zinc, will be considered.

Appearance

Most of the d-block elements are considered to be metals, with a common lustrous metallicappearance. All but copper and gold have a silvery colour.

General Reactivity

These elements have d electrons in their valence shells, and this gives them differentcharacteristics to other metals in the periodic table. They each exist in several oxidation statesexcept scandium and zinc; many of their compounds are coloured; and they readily form complexes

by acting as Lewis acids.

Occurrence and Extraction

The first six elements, scandium to iron, occur mainly as the oxides in various mineral deposits.The most abundant of these is iron, found chiefly in magnetite and haematite, both commonlyknown as iron ore. The remainder of the elements occur mainly as sulfides such as zinc blende.

Each element is extracted from the appropriate mineral by various extraction methods. Theextraction of iron, however, is of immense importance as steel - basically a mixture of iron andcarbon - is used in greater quantities world-wide than any other metal. Steel is produced from ironore in two main stages:

(1) a blast furnace produces impure iron from iron ore(2) the impure iron is then purified and alloyed with other metals to produce steel.

Physical Properties

All these elements are hard, rigid and have good thermal and electrical conductivities. They havehigh melting and boiling points.

Chemical Properties

The chemistry of the d-block elements is governed by the fact that most exhibit several oxidationnumbers. This is because the energies of all the d electrons are very similar. The d electrons alsoconfer properties on these elements not found elsewhere:- they easily form complexes

- their complexes are often coloured- some complexes are paramagnetic- they make good catalysts.

The chemical properties of these elements and their many complexes is extensive, and not

suitable for further study here.

Oxidation States


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As stated earlier, most of the d-block elements exist in several oxidation states - for example, theoxidation number of iron can be 0, +2, +3 and +6. The widest range of oxidation numbers is formanganese, which has a lowest oxidation number of 0 and a highest oxidation number of +7. Thereare general tendencies concerning the oxidation numbers:

(a) the 1st and last elements, scandium and zinc, have only 1 oxidation number.(b) all the elements except zinc can have oxidation number +3(c) all the elements except scandium can have oxidation number +2

(d) from scandium to manganese, the highest oxidation number = the number of 4s electrons +the number of 3d electrons

(e) from manganese to zinc, low oxidation numbers are common.

Industrial Information

The d-block elements are used in many thousands of applications.Iron is the most widely used element because it is converted to steel, which consists of iron with

0.2 - 1.7% carbon. The addition of carbon hardens the iron and gives it better resistance tocorrosion. Special steels can be prepared by the addition of small quantities of other elements -stainless steel contains 18% chromium and 8% nickel. Iron and steel are extensively used in oursociety.

Other important uses of some of these elements include titanium in aircraft and spaceshipmanufacture. Titanium is less dense than other d-block elements, and this lightness, coupled with itsextra hardness, make it more suitable than aluminium in high-flying aircraft and space vessels.

Chromium is often used for electroplating, and alloyed with nickel to make nichrome - used inelectrical components as its electrical resistance hardly varies with temperature.

Copper is used as protective sheeting as it is more resistant to oxidation than other elements. Thegreen patina that forms on exposure to air is copper(II) carbonate, sulfate or chloride. Copper is alsoused in electrical cables.

Zinc is used in protecting steel by galvanising, and in alloys, eg brass (copper and zinc).

Variable oxidation state

The multiple oxidation states of the d-block (transition metal) elements are due to the proximityof the 4s and 3d sub shells (in terms of energy). All transition metals exhibit a 2+ oxidation state(both electrons being lost from the 4s and all have other oxidation states (common). ( Table 4.1)

Sc Ti V Cr Mn Fe Co Ni Cu Zn

+1

+2 +2 +2 +2 +2 +2 +2 +2 +2

+3 +3 +3 +3 +3 +3 +3 +3+4 +4 +4 +4

+5

+6 +6 +6+7

Table 4.2

Coordinated ligands

Ligands are the molecules (or ions) which donate an electron pair to form a dative covalent bondwith the central transition metal atom (forming a complex molecule or ion).

Coordination Complexes

These are species which are formed around a central atom, with other atoms, ions or moleculesdonating an electron pair to form a covalent bond to this central atom. The result is a "complex"usually an ion but may also be a molecule. (Table 4.2)

Complex Shape Ligands Coordina-

tion number Name


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[Fe(H2O)6]3+ octahedral Water 6 hexa-aqua iron III ion

[Fe(CN)6]3- octahedral cyanide CN- 6 hexacyano ferrate III ion

[CuCl4]3- tetrahedral chloride Cl- 4 tetrachloro cuprate I ion

[Cu(NH3)4]2+ square planar Ammonia 4 tetra-ammine copper II ion

[Ag(NH3)2]+ linear Ammonia 2 diammine silver I ion

Ni(CO)4 tetrahedral carbon

monoxide 4

tetracarbonyl nickel 0molecule

Table 4.3

Coloured compounds

The color in the transition metals (d-block) is predominantly due to the splitting of the d shellorbitals into slightly different energy levels. As a result, certain wavelengths of energy can be

absorbed by the d-block elements (with electrons jumping between these slightly different energylevels), resulting in the complement color being visible.

Colour is affected by both the oxidation state of the transition metal and the type of ligand(Table 4.3)

Complex ion Oxidation state of

metal Colour Ligand

[Fe(H2O)6]3+ III pale green Water

[Fe(H2O)6]2+ II yellow Water

[Cu(H2O)6]2+ II blue Water

[Cu(NH3)4]2+ II deep blue ammonia

[CuCl4]2 II green chloride ion

Table 4.4

Crystal field theory

Magnetism

Transition metals and their ions often have unpaired 'd' electrons which produce an asymmetric

magnetic field that can be detected. This is called paramagnetism (Table 4.4)

Examples:

Complex ion Electronic

configuration

No of unpaired

electrons Magnetism

[Fe(H2O)6]3+ [Ar]4s0 3d5 5 paramagnetic

[Cr(H2O)6]3+ [Ar]4s0 3d3 3 paramagnetic

[Cu(H2O)6]2+ [Ar]4s0 3d9 1 paramagnetic

[Ni(NH3)6]2+ [Ar]4s0 3d9 2 paramagnetic

[CoCl4]2- [Ar]4s0 3d7 3 paramagnetic


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lectures part 1

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