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notes -periodic properties

Long Form of the Periodic Table or Modern Periodic Table ;

The modern periodic table is based on modern periodic law given by Moseley in 1942. Themodern periodic law states that the physical and chemical properties of the elements areperiodic function of their atomic numbers. Thus, when the elements were arranged in theorder of their increasing atomic numbers, the elements of similar properties recur at regularintervals.

This table consists of horizontal rows called as periods and vertical columns called asgroups.

Periods ; There are seven periods in the periodic table and each period starts with a differentprincipal quantum number.

The first period corresponding to n = 1 consists of only two elements hydrogen (1s1) and

helium (1s2). This is because the first energy shell has only one orbital (1s), which canaccommodate only two electrons.

In the second period corresponding to n = 2, there are four orbitals (one 2s and three 2p)having a capacity of eight electrons and so contains eight elements. This period starts withlithium (Z = 3) with electron entering the 2s orbital and ends with neon (Z = 10) where the

second shell is complete (2s2 2p6).

In the third period corresponding to n = 3, there are nine orbitals (one 2s three 2p and five3d). As 3d orbitals are higher in energy, they are filled after the 4s orbitals.

The fourth period corresponding to n = 4, consists of filling of one 4s and three 4p orbitals..there are eighteen elements in this period starting from potassium with electron entering the 4s

orbital (Z = 19) to krypton (Z = 36) where the third shell gets completed (4s2 3d104p6).

In the fifth period there are 18 elements like the fourth period. It begins with rubidium (Z = 37)with the filling of 5s orbital and ends with xenon (Z = 54) with the filling of 5p orbital.

The sixth period contains 32 elements (Z=55 to 86) and the successive electrons enter into 6s,4f, 5d and 6p orbitals in that order. It starts with caesium and ends with radon.

The seventh period, though expected to have 32 elements is incomplete and contains only 19elements at present. Number of elements in each period ; There is a periodicity occurring atregular intervals of 2, 8, 8, 18, 18 and 32 and so the numbers 2, 8, 18 and 32 are called magicnumbers. The first three periods are called short periods while the other three periods are calledlong periods.

Groups; The vertical column in the periodic table is called as group. There are 18 groups in thelong form of the periodic table and they are numbered from 1 to 18 in the IUPAC system.

Merits of the long form of the periodic table

1. based on the most fundamental property of the elements the atomic number, so it is moreaccurate.

2. the position of isotopes in one place is justified.

3. four blocks of s, p, d and f has made the study of elements simpler.

4. The position of the elements, which were misfit on the basis of atomic mass is now justifiedon the basis of atomic number

Grouping of Elements In the long form of the periodic table,; elements are grouped into fourmain blocks, purely on the basis of electronic configurations. Elements are grouped in blocks s,p, d and f depending on the nature of orbital(s) into which the last electron of the atom enters.

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The s block elements ; s block elements, (except helium) are those in which the valence-shellis the ns orbital and where the last electron enters into the s orbital of the outer element. The

general outer electronic configuration of such elements is either ns1or ns2. The s blockelements belong to group 1 and 2 of the long form periodic table. They are situated on the leftside of the table.

The p Block Elements p block elements are those in which the outer electronic configuration isof the type

ns2 np1to6 and where the last electron enters into any of the outermost p orbitals. Theseelements belong to groups 13 to 18 of the long form of periodic table and are situated on the righthand side of the table.

The d Block Elements d block elements are those in which the added electron goes into one ofthe d orbitals. These elements have valence electrons in both their outermost and penultimate

shells (second outermost) and have a general outer electronic configuration of (n 1) d1-10 ns1-

2. The penultimate shell in these elements is expanded from 8 to 18 by the inclusion of ten delectrons. It is for this reason that these elements are called d block elements. Thus, d blockelements are those, which in their elemental or combined forms have partially filled d orbitals.They are also referred to as transition elements. d block elements belong to groups 3 to 12 ofthe long form of periodic table and are situated in the middle of the table between s block and pblock elements. All d block elements are classified into four transition series, namely: 3d

series in the 4th period having 10 elements 4d series in the 5th period having 10 elements 5d

series in the 6th period having 10 elements 6d series in the 7th period having incompleteelements.

The f Block Elements The f block elements are those, which in their elemental or ionic formshave partially filled f orbitals. The differentiating (last) electron enters f orbitals, which lie inner tothe second outermost (penultimate) shell. Thus these elements are also known as inner-transition elements. There are two series of f block elements, each having 14 elements.Lanthanides (atomic number 58 71) are those inner-transition elements in which 4 f orbitalsare progressively filled. Actinides (atomic number 90 103) are those elements in which 5 forbitals are progressively filled. The f block elements are placed at the bottom of the long formof periodic table in the form of two rows. Most actinides are radioactive. The general outer

electronic configuration of f block elements is, (n 2) f1-14 (n 1) d0-1 ns2

Representative or Main Group Elements; These consist of all s and p block elementsexcluding the noble gases (group 18 elements). The chemical properties of the representativeelements are determined by the number of valence electrons in their atoms. The number ofvalence electrons belonging to the above two blocks are: s block elements: 1 or 2: Same as thegroup number p block elements: 3 to 8: Group number minus ten The representativeelements are also known as normal or typical elements. Thus on the basis of their location in theperiodic table, these are elements, which belong to groups 1, 2 and 13 to 17.

Noble gases or Aerogens These consist of gaseous elements of group 18, the last group in pblock. All noble gases have a general outermost electronic configuration of ns and np. The only

exception is helium having an outermost configuration of 1s2. Because they display these stablecharacteristics (having completely filled outermost shells), they do not show any chemicalreactivity. Only Xenon shows limited chemical reactivity under severe conditions.

Transition elements d block elements are called transition elements. These elements occupygroups 3 to 12 on the periodic table. There are four series of transitional elements, 3d, 4d, 5dand 6d depending on the energy levels of d orbitals. They have a general outermost electronic

configuration of (n 1) d1-10 ns1-2.

Inner transition elements The f block elements are called inner transition elements. There aretwo series of f block elements. These are the 4f and 5f series called Lanthanides and

Actinides. The general outer electronic configuration of f block elements is, (n 2) f1-14 (n 1)

d0-1 ns2.

Periodic Properties;

Atomic radii ; defined as one half of the inter-nuclear distance between the two centres of theneighbouring atoms in a covalently bonded molecule.

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Variation of Atomic Radii

Variation in a period Atomic radii in general, decrease with increase in atomic number, goingfrom left to right in a period. This is explained on the basis of increasing nuclear charge along aperiod. The nuclear charge increases progressively by one unit while the corresponding additionof one electron takes place in the same principal shell. As the electrons in the same shell do notscreen each other from the nucleus, the nuclear charge is not neutralized by the extra valenceelectron. Consequently the electrons are pulled closer to the nucleus by the increased effectivenuclear charge resulting in the decrease in the size of the atom. In this way the atomic size goeson decreasing across the period.

element Li Be B C N O F

Atomic radius-pm 135 90 82 70 70 66 64

The atomic radius abruptly increases in the case of noble gas element Neon as it does not formcovalent bonds. So the value of Neon radius is Van der Waals radius which is considerablyhigher than the value of other covalent radii.

Variation in a group The atomic radii of elements increases from top to bottom in a groupbecause the nuclear charge increases with increasing atomic number. Although, there is anincrease in the principal quantum number from one atom to another, the number of electrons inthe valence shell remain the same. The effect of increase in the size of the electron cloud outweighs the effect of increased nuclear charge and so the distance of the valence electron fromthe nucleus increases down the group. Thus the size of the atom goes on increasing down thegroup in spite of increasing nuclear charge.

element Li Na K Rb Cs

Atomic radius 135 154 196 211 225

Ionic Radii These are radii of ions in ionic crystals. Ionic radius may be defined as the effectivedistance from the center of nucleus of an ion up to which it has an influence on its electron cloud.In ionic compounds the inter nuclear distance may be taken as equal to the sum of the ionic radiiof the two ions. The inter nuclear distance in ionic crystals are obtained from X-ray studies.

Radius of the cation A cation is formed by the loss of one or more electrons from the gaseousatom. Thus, the whole of the outer most shell of electrons is removed resulting in the smallersize in the cation. For example, in lithium atom, there is only one electron in the outermost 2s

shell. As the lithium atom changes to Li+ ion the outer most 2s shell disappears completely.This disappearance results in the decrease in size.

With the removal of electrons from an atom the magnitude of the nuclear charge remains thesame while the number of electrons decreases. As a result the nuclear charge acts on lessnumber of electrons. The effective nuclear charge per electron increases and the electrons aremore strongly attracted and pulled towards the nucleus. This causes a decrease in the size of the



ion.

Radius of the anion The negative atom is formed by the gain of one or more electrons in theneutral atom. The number of electrons increases while the magnitude of nuclear charge remainsthe same. The same nuclear charge acts on larger number of electrons than were present in theneutral atom. The effective nuclear charge per electron is reduced and the electron cloud is heldless tightly by the nucleus. This causes an increase in the size of the ion. Thus anions are largerin size than the corresponding atom.

Variation of ionic radii in a group The ionic radii in a particular group increases in moving fromtop to bottom because of the increase in the principal quantum number though the number ofelectrons in the valence shell remains the same.

Isoelectronic ions are ions with same number of electrons.

In isoelectronic series of ions, as the nuclear charge increases the electrons are pulled moreand more strongly and the size decreases.

ion Na + Mg 2+ Al 3+ P 3- S 2- Cl -

Ionic radius 95 65 50 212 184 181

Ionization Energy ;The amount of energy required to remove the most loosely bound electronfrom an isolated gaseous atom is called ionization energy (IE).

Ionization energy is also called as ionization potential because it is measured as theminimum potential required to remove the most loosely held electron from the rest of theatom. It is measured in the units of electron volts (eV) per atom or kilo joules per mole of

atoms (kJ mol-1). 1 eV per atom = 96.64 kJ mol-1 = 23.05 k cal mol-1 Thus, the ionizationenergy gives the ease with which the electron can be removed from an atom. The smaller thevalue of the ionization energy, the easier it is to remove the electron from the atom.

.

Factors affecting Ionization Energy; 1. Size of the atom 2.Charge on the nucleus 3.Screeningeffect . 4.Electronic arrangement

Variation along a period ; The ionization energy increases with increasing atomic number ina period. This is because 1) The nuclear charge increases on moving across a period from leftto right. 2) The atomic size decreases along a period though the main energy level remains thesame. Due to the increased nuclear charge and simultaneous decrease in atomic size, thevalence electrons are more tightly held by the nucleus. Therefore more energy is needed toremove the electron and hence ionization energy keeps increasing.

However some irregularities have been noticed due to the extra stability of the half filled andcompletely filled configurations. For example, the nuclear charge on Boron is more thanBeryllium, yet there is slight decrease in ionization energy from Be to B. This is because, in boronthe last electron goes to 2p orbital which is at a slightly higher energy than 2s orbital. Also, theelectronic configuration of B is less stable than that of Be (has completely filled orbitals). Hencethe ionization energy is less than that of Be. Similarly, nitrogen, which has half filled 2p orbitals,is more stable than oxygen. Therefore the ionization energy of nitrogen is more than that ofoxygen.

element Li Be B C N O F Ne

Ionizationenergy

520 899 801 1056 1402 1314 1651 2050

Variation down a group ; The ionization energy gradually decreases in moving from top tobottom in a group. This is due to the fact that: 1)The nuclear charge increases in going from topto bottom in a group. 2)An increase in the atomic size due to an additional energy shell (level)n. 3)Due to the increase in the number of inner electrons there is an increase in the shieldingeffect on the outer most electron. The effect of increase in atomic size and the shielding effect is



much more than the effect of increased nuclear charge. As a result , the electron becomes lessfirmly held to the nucleus and so the ionization energy decreases as we move down the group.


Ionization energy 520 496 419 403 376

Electron Affinity; Electron affinity is the amount of energy released when an electron is added toan isolated gaseous atom.

Electron affinity is the ability of an atom to hold an additional electron. If the atom has moretendency to accept an electron then the energy released will be large and consequently theelectron affinity will be high. Electron affinities can be positive or negative. It is taken as positivewhen an electron is added to an atom. It is expressed as electron volts per atom (eV per atom) orkilo joules per mole.

Variation along a period - The size of an atom decreases and the nuclear charge increases onmoving across a period. This results in greater attraction for the incoming electron. Hence theelectron affinity increases in a period from left to right.

Variation down a group As we move down a group the atomic size and nuclear size increases.As the effect of increase in atomic size is more pronounced the additional electron feels lessattracted by the large atom. Consequently the electron affinity decreases.

Electronegativity The relative tendency of an atom in a molecule to attract a shared pair ofelectrons towards itself is termed as Electronegativity.

The value of electronegativity of an element describes the ability of its atom to compete forelectrons with the other atom to which it is bonded. Electronegativity is however not the property ofan isolated atom. Electronegativity is measured on a number of scale levels, the most commonlyused are of Pauling or Mulliken.

Factors Affecting Electronegativity

1)Atomic size 2)Ionisation energy and electron affinity 3)Number and nature of atoms 4)Typeof hybridization 5)Charge on the ion

Variation along a period - As the nuclear charge increases from going left to right in a periodbecause the electrons enter the same shell, the shielding is less effective. Thus the increasednuclear charge attract the shared pair of electrons more strongly resulting in higherelectronegativity from going left to right in a period.

element Li Be B C N O F

electronegetivity 1.0 1.5 2 2.5 3 3.5 4

Variation down the group - Electronegativity decreases down the group because the atomic sizeincreases. The larger the size of the atom the lesser the tendency to attract the shared pair ofelectrons.


electronegetivity 1 0.9 0.8 0.8 0.7

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