Periodic table and properties-final.pptx.pdf

Periodic table and properties Periodic Table of Elements Periodic table was developed to avoid the difficulty of studying the properties of all the elements individually. Studying the properties of groups or families of elements those have similar properties. The first importan...

Periodic table and properties Periodic Table of Elements Periodic table was developed to avoid the difficulty of studying the properties of all the elements individually. Studying the properties of groups or families of elements those have similar properties. The first important attempt to divide elements based on their physical and chemical properties was done by the Russian chemist Dimitri Mendeleev in 1869. Mendeleev’s periodic law states that“Physical and chemical properties of elements are periodic functions of their atomic weights”. Periodic Table of Elements In 1912 British chemist, Henry Mosely introduced the modern periodic law which states that “ The properties of elements are periodic functions of atomic numbers”. Periodic Table of Elements Periodic table is a table in which elements are arranged in the Modern form of the periodic table order of increasing atomic long form of periodic table number in the manner that the elements with similar properties fall in the same vertical column. Elements which fall in the same vertical column are called group of elements. They resemble closely with one another in their properties. Periodic Table of Elements Long form or modern form of periodic Table Long form of periodic table helps in organizing and systematizing the chemistry of the elements. No of elements known at present is 118. Long form of periodic table also helps us to understand the reason why certain elements resembles one another and why they differ from other elements. It helps us to understand the periodicity of properties and why some properties recur in regular periodic intervals. Periodic Table of Elements Long form or modern form of periodic Table Periods : These are the horizontal rows of periodic table. There are seven periods in periodic table. First period –Hydrogen and Helium Second period- 8 elements (Li-Ne) Third period- 8 elements (Na-Ar) Fourth period- 18 elements (K-Kr). K, Ca, Ga, Ge, As, Se, Br and Kr are main group elements. Sc to Zn are called transition elements. First long period. Fifth period- 18 elements (Rb-Xe). Sixth period -32 elements(Cs-Rn), 8normal elements, 10 transition elements 14 inner transition elements. Seventh period-(Fr onwards )32 elements Actinides belong to this group. Periodic Table of Elements Lanthanides and actinides Lanthanides (At no 57-71) belong to sixth period. They are also called f-block elements and also called rare earth elements. They are placed separately below the periodic table to avoid undue expansion of table. Actinides (At. No 93-112) are the fourteen elements after actinium. The elements after uranium are called transuranic elements and they are prepared artificially. Periodic Table of Elements Groups Elements arranged in the vertical column of the periodic table are called groups. There are 18 groups in the periodic table. Group 1- Hydogen and alkali metals Group 2- Alkaline earth elements Group 3-III B- Sc, Y, Lanthanides , Actinides Group 4- IV B- Ti, Zr, Hf, Rf Group 5- V B- V, Nb, Ta Group 6- VI B- Cr, Mo, W Group 7- VII B- Mn, Tc, Re Group 8- VIII B- Fe, Ru, Os Group 9- VIII B- Co, Rh, Ir Group 10- VIII B- Ni, Pd, Pt Group 11- I B- Cu, Ag, Au Group 12- II B- Zn, Cd ,Hg Periodic Table of Elements Groups Group 13- III A- Boron group Group 14- IV A- Carbon group Group 15- VA- Nitrogen group Group 16- VI A- oxygen group Group 17- VII A- Halogen group Group 18- VIII A- Noble gas group Numbering according to roman numerals are based on the no of electrons in the outermost shell. Periodic Table of Elements Periodicity: the recurrence of similar properties of elements after a regular interval when they are arranged in the increasing order of atomic numbers. Cause of periodicity: Similar outer electronic configuration. Periodic Table of Elements Division of elements in the periodic table Elements of periodic table are divided into 4 groups i.e. S, p, d and f elements These classifications are based on the orbital into which last electron of the atom of the element enters. s block elements - Those elements of the periodic table in which last electron enters the s orbital of the outer most shell are called s block elements. Since s orbital can accommodate only two electrons there are only two groups in this category. General outer electronic configuration ns1-2 Alkali (ns1)and alkaline earth elements (ns2) p block elements - Those elements of the periodic table in which last electron enters the p orbital of the outer most shell are called p block elements. Since p orbital can accommodate only six electrons there are six groups in this category. General outer electronic configuration ns2, np1-6 2 1 2 2 Group 13-boron 2 3 group(ns np ), group 14-carbon 2 4 group(ns np ), group 15-nitrogen 2 5 group(ns np ), group 16-oxygen2 6 group(ns np ), group 17-halogen group(ns np ) and group 18-inert gas group(ns np ) Periodic Table of Elements Division of elements in the periodic table d block elements - Those elements of the periodic table in which last electron enters the d orbital are called d block elements. The last electron enters the d orbital which is the penultimate shell. Since d orbital can accommodate 10 electrons there 10 groups in this category. General electronic configuration – (n-1)d1-10 ,ns2 Sc- 3d1, 4s2, Y-4d1, 5s2 There are fourseries of transition elements First series –Sc to Zn, Second series- Y to Cd, the third series – La –Hg and the fourth series- Ac-Cn F block elements - Those elements of the periodic table in which last electron enters the f orbital are called f block elements. Last electron enters the f orbital whih is the pre-penultimate shell. Since f orbital can accommodate 14 electrons there are 14 elements in this category. General electronic configuration –(n-2)f1-14, (n-1)d1 ,ns2 Ce- 4f1,5d1, 6s2 Eu- 4f7, 5d1, 1, 6s2 There are two series of f block elements First series – Lanthanides -General electronic configuration [Xe] 4f1-14 5d1 6s2 Ce- 4f1,5d1, 6s2 Eu- 4f7, 5d0, 6s2 Lu- 4f14,5d1, 6s2 Second series- Actinide series- General electronic configuration [Rn] 5f1-14 6d1 7s2 Th- 5f1,6d1, 7s2 U- 5f3, 6d1, 7s2 Periodic Table of Elements Division of elements in the periodic table Write the electronic configuration of Uranium Write down the electronic conjuration of Cesium Write the electronic configuration of fluorine Write down the electronic conjuration of Germanium Write the electronic configuration of Gallium Write down the electronic conjuration of Molybdenum Write down the electronic configuration of the element which belongs to 14 th group and 6th period Periodic Table of Elements Atomic properties Sizes of atoms and ions- atomic and ionic radii Atomic radii Atomic radius is the distance between the center of the nucleus and outer most shell of an atom. In period atomic radii decreases on moving from left to right. In a period on moving from left to right nuclear charge increases and electrons are added to the same shell. So they are attracted more strongly towards nucleus and size decreases. Periodic Table of Elements Atomic properties Atomic radii In a group on moving from top to bottom atomic size increases. On moving from top to bottom no of shell increases and as result size of atom increases. Periodic Table of Elements Atomic properties Ionic radii Ionic radius is the radius of ion in an ionic crystal. Ionic radius is the distance between the centre of nucleus to the distance where the nuclear charge has influence on electron cloud. In the case of cations, ionic radii is smaller than atomic radii. In the formation of cations in the case of alkali and alkaline earth elements the removal of electrons causes the removal of the outermost shell. This causes a decrease in ionic radii. In the formation of anions, one or more electrons are added to the outermost shell but effective nuclear charge remains the same. So the effective nuclear charge and attractive force experienced by the electrons decreases and size increases Periodic Table of Elements Atomic properties Ionic radii Periodic Table of Elements Atomic properties Ionisation energy Ionisation energy (IE) is the energy required to release the outermost shell of an isolated gaseous atom. M + IE → M+ + e- Elements with lower ionization potential will form ions rapidly. Inoisation energy is usually expressed in electron volt(eV). It may also be expressed in Joules or kilojoules. Energy required to release the first electron from outermost shell is called first ionization energy (IE1) and energy required to release the second electron from outermost shell is called second ionization energy (IE2) and further IE3, IE4 etc. IE1< IE2 < IE3 < IE4 When one electron is removed from the atom, the effective nuclear charge on the remaining electrons increases and they are held closer to the nucleus and more energy needs to be applied to release the next electron from a mono-positive ion. Periodic Table of Elements Atomic properties Ionisation energy Factors affecting ionization energy 1. Atomic size 2. Nuclear charge 3. Number of electrons in the inner shells 4. Penetration effect or the effect of removal of s, p, d and f electrons 5. Electronic configuration Periodic Table of Elements Atomic properties - Ionisation energy Factors affecting ionization energy 1. Atomic size Larger the atomic size, smaller is the ionization energy As the size of the atom increases, outermost electrons are farther away from the nucleus According to coulomb’s Law attractive pull of nucleus on outermost electron decreases Less energy is required to knock out the electron Periodic Table of Elements Atomic properties -Ionisation energy Factors affecting ionization energy 2. Nuclear charge The force of attraction between the nucleus and electron increases with nuclear charge. As the nuclear charge increases more energy is required to release the electron from the atom. Hence ionization energy increases with increase in nuclear charge. Periodic Table of Elements Atomic properties -Ionisation energy Factors affecting ionization energy 3. Number of electrons in the inner shells As the number of electrons in the inner shell increases ionization energy decreases. The electrons in the inner shell acts as a shield or screen between nucleus and electrons in the outermost shell. It is called shielding effect. As the number of electrons in the inner orbitals increases shielding effect increases and effective atomic number on the outermost electron decreases. Ionisation energy, energy required to release the electrons from the outermost shell decreases. Periodic Table of Elements Atomic properties Ionisation energy Factors affecting ionization energy 4. Penetration effect or the effect of removal of s, p, d and f electrons As s electrons move around the nucleus, it has more probability of coming closer to the nucleus than p. d and f electrons of the same principal shell. S electrons penetrate more towards nucleus than p, d and f electrons and the penetration power of electrons in the same shell s>p>d>f. Thus s electron experience more attracting power from nucleus compared to s, p, d and f electrons So energy require to pull out s electron is maximum and f electrons the least. Ionisation energy corresponding to the elimination of an electron from a given energy level varies in the order s>p>d>f. Periodic Table of Elements Atomic properties -Ionisation energy Factors affecting ionization energy 5. Electronic configuration Certain electronic configuration is stabler than others. Half-filled and full-filled electronic configurations are stabler compared to other configuration. So it is required to apply more energy to release electrons from the stabler configuration. Ionisation energy is higher for half-filled and fully-filled electronic configurations. Periodic Table of Elements Atomic properties -Ionisation energy Variation of ionisation energy in the periodic table Variation in the period Ionisation energy increases in a group moving from left to right In a group moving from left to right Atomic size decreases Effective nuclear charge increases The ionization energy becomes maximum for inert gases which has a stable electronic configuration Periodic Table of Elements Atomic properties - Ionisation energy Variation of ionisation energy in the periodic table Variation in the group In general ionization energy decreases on moving from top to bottom in a group. Atomic size increases Effective attraction of nucleus on outermost charge decreases. Screening effect increases. Periodic Table of Elements Atomic properties -Ionisation energy Variation of ionisation energy in the periodic table Periodic Table of Elements Atomic properties Ionisation energy Periodic Table of Elements Atomic properties – Electron affinity Electron affinity is the amount of energy released when an electron is added to a gaseous isolated atom or ion. X(g) + e- → X- + Energy (EA) Greater the energy released while taking up an extra electron greater is the electron affinity. Electron affinity of the atom measures the tightness with which an atom binds an additional electron to itself. It is expressed in eV/atom or kJ/mol It is a measure of ease with which an anion is formed. Greater the ease with which anion is formed, greater is the electron affinity , greater is the energy released. Periodic Table of Elements Atomic properties – Electron affinity Electron affinity is the amount of energy released So it is an exothermic reaction The energy released has a negative sign. Periodic Table of Elements Atomic properties – Electron affinity Variation of electron affinity in the periodic table In general electron affinity decreases in going from top to bottom in a group and increases in going from left to right across the period. On going down the group atomic size increases and attraction of nuclear charge on outer electron decreases. Atoms will have less tendency to attract additional electrons to themselves. Electron affinity decreases. On moving across the period, atomic size decreases and effective nuclear charge increases. Atoms can easily accommodate additional electrons to themselves. Electron affinity increases. Electron affinities of metal are less and non-metals are more. Periodic Table of Elements Atomic properties – Electron affinity Variation of electron affinity in the periodic table Halogens have maximum electron affinity. This is due to their strong tendency to gain an electron to acquire stable electronic configuration. Electron affinity decreases in a group. In the case of halogens, from Cl to I electron affinity decreases. But Florine has an exceptionally low electron affinity. It is probably due to the small size of the atom. The addition of an electron produces high electron charge density in a relatively compact 2p subshell resulting in strong electron-electron repulsion. The repulsive force between electrons implies low electron density. Electron affinities of inert gases are zero. They have a stable s2p6 electronic configuration and no chance of the addition of an electron to it. Some elements have practically zero electron affinity. This is because of half-filled and full-filled electronic configurations. Eg Be, Mg, Ca, and N. Periodic Table of Elements Atomic properties – Electron affinity Variation of electron affinity in the periodic table Periodic Table of Elements Atomic properties Electronegativity Electronegativity is the power of an atom in a molecule to attract an electron towards it. The ability of an atom in a molecule to attract an electron towards it depends upon its environment in the molecule. Electronegativity of an atom in a molecule therefore depends on the other atoms present in the molecule. Periodic Table of Elements Atomic properties Electronegativity Factors influencing the electronegativity 1. Charge on the atom 2. Hybridisation 3. Effect of substituents 4. Role of ionization energies and electron affinities 5. Effective nuclear charge Periodic Table of Elements Atomic properties Electronegativity Factors influencing the electronegativity 1. Charge on the atom Atoms with positive charge (Cations) in a molecule will be more electronegative than anions. As the positive charge increases electronegativity increases Anions will have less electronegativity The greater the positive power of an atom in a molecule greater will be its affinity to attract electrons. Atoms in a molecule with more positive oxidation state will have more electronegativity. Eg. HClO and HClO3 In HClO oxidation state of chlorine is +1 In HClO3 oxidation state of chlorine is +5. So chlorine in HClO3 is more electronegative Periodic Table of Elements Atomic properties Electronegativity Factors influencing the electronegativity 2. Hybridisation The penetrating power of electrons increases in the order S>P>d>f More nuclear attraction is experienced by s electrons. In molecules with sp hybridization where 50% s character will experience more nuclear power on it. So atoms in these molecules experience more nuclear charge, hence it can easily attract an electron to it. So electronegativity will be more for atoms in an sp hybridized molecule Carbon in acetylene(sp hybridization) will be more electronegative than carbon in ethylene (sp2 hybridization) and ethane (sp3 hybridization) Periodic Table of Elements Atomic properties Electronegativity Factors influencing the electronegativity 3. Effect of substituents The electronegativity depends on the substituents on a molecule E.g. CH4 and CF4 Carbon in CF4 is more positive than carbon in CH4. So carbon in CF4 is more electronegative Periodic Table of Elements Atomic properties Electronegativity Factors influencing the electronegativity 4. Role of ionization energies and electron affinities Electronegativity is the average of electron affinity and ionization energy. So higher the electron affinity and ionization energy, the higher is the electronegativity Periodic Table of Elements Atomic properties Electronegativity Factors influencing the electronegativity 5. Effective nuclear charge 1. Electron attracting power of an atom in a molecule is proportional to the effective nuclear charge. So any factor which increases the nuclear charge will increase electronegativity. Periodic Table of Elements Atomic properties Electronegativity Variation of electronegativity in the periodic table 1. Electronegativity decreases on going down the group and increases on moving across the period. Periodic Table of Elements Atomic properties Electronegativity Variation of electronegativity in the periodic table Element with the lowest electronegativity- Francium Element with the higest electronegativity- Fluorine Periodic Table of Elements Atomic properties Metallic characcter Variation of metallic character in the periodic table Periodic Table of Elements  Characteristic properties of s-block elements They have low-density, melting and boiling points They are good conductors of electricity They are metals. Lower ionization potential Low electronegativity. They form ionic bonds with a fixed oxidation state which is equal to the number of outermost electrons They impart characteristic colors to the flame. They have weak tendency to form complexes. The first two atoms of both groups differ from rest in their properties.  Outer most electrons of the ground atoms absorb energy from flame and get exited to the higher energy level. This state is referred as exited state. Exited states are very unstable and short lived states. Atom comes down to ground state by releasing extra energy in the form of radiation. This radiation for alkali metals is in the visible region and hence they appear as different colors.  P-block elements The elements belonging to groups 13-18 constitute p block elements. They are solids/liquids/gases at room temperature (Br is liquid) They form acidic oxides They impart no characteristic color to the flame They form covalent compounds. Halogens form salts with alkali metals They have high ionization potentials They have very large electron gain enthalpies The aqueous solutions their oxides are acidic in nature d-block elements –Transition elements Four periods of 10 elements that are placed between s-block and p-block elements. They are also called d-block elements because the last electron enters the d orbitals last but one (penultimate)energy shell. They are called transition elements because these represent a complete change from most electropositive s-block elements to the least electropositive p-block elements. There are four series of transition elements First series Sc-Zn second series Y-Cd Third series La- Hg Fourth series Ac -Cn Electronic configuration of elements of first series of transition elements Ar At No 18 Electronic configuration 1s2 2s2 2p6 3s2 3p6 Cr and Cu shows unexpected electronic configuration due to the stability of half filled and completely filled orbitals Characteristics of transition elements They show variable valency They form coloured compounds They show paramagnetic properties They forms coordination complexes Chemical and physical properties of transition elements are governed by electronic configuration the number of electrons in d orbitals.  Lanthanides Lanthanides Lanthanum (z=57) and the next fourteen elements (Z=58 - 71) which follow it are called lanthanides. They are also called f block elements because they have partially filled f orbitals and the last electron enters f orbital (Pre-penultimate). There are fifteen lanthanides because f orbitals can accommodate 14 electrons. These fifteen elements resemble one another very closely. They have similar chemical properties. They have similar outer electronic configuration. They differ only in 4f electrons which are embedded interior to 5d and 6s electrons which are alike are exposed to surroundings. All elements are naturally occurring except promethium. Electronic configuration of lanthanides Element Symbol Electronic configuration Electronic configuration of Xe Lanthanum La [Xe] 4f0 5d1 6s2 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 4d10 5s2 5p6 Cerium Ce [Xe] 4f1 5d1 6s2 ([Xe] 4f0 5d2 6s2 ) 1-14 1 2 Praseodymium Pr [Xe] 4f2 5d1 6s2 General electronic configuration [Xe] 4f 5d 6s Neodymium Nd [Xe] 4f3 5d1 6s2 Promethium Pm [Xe] 4f4 5d1 6s2 Samarium Sm [Xe] 4f5 5d1 6s2 Europium Eu [Xe] 4f7 6s2 Gadolinium Gd [Xe] 4f7 5d1 6s2 Terbium Tb [Xe] 4f8 5d1 6s2 Dysprosium Dy [Xe] 4f9 5d1 6s2 Holmium Ho [Xe] 4f10 5d1 6s2 Erbium Er [Xe] 4f11 5d1 6s2 Thulium Tm [Xe] 4f12 5d1 6s2 Ytterbium Yb [Xe] 4f14 6s2 Lutetium Lu [Xe] 4f14 5d1 6s2  Atomic Name Symbol Electronic Oxidation states Number configuration 57 Lanthanum La [Xe]4f0 5d1 6s2 +3 58 Cerium Ce [Xe]4f05d2 6s2 +3, +4 1. Like transition metal they do not show variable valency. 59 Praseodymium Pr [Xe]4f2 5d1 6s2 +3 2. The common oxidation state of all the 60 Neodymium Nd [Xe]4f35d1 6s2 +3 lanthanoids is +3. This state is attained by losing 5d and 6s electrons. 61 Promethium Pm [Xe]4f4 5d1 6s2 +3 3. Lanthanides form colored complexes due to 62 Samarium Sm [Xe]4f5 5d1 6s2 +3 f-f transitions. 63 Europium Eu [Xe]4f7 5d0 6s2 +3, +2 4. They also show paramagnetism. 64 Gadolinium Gd [Xe]4f7 5d1 6s2 +3 65 Terbium Tb [Xe]4f7 5d2 6s2 +3, +4 66 Dysprosium Dy [Xe]4f9 5d1 6s2 +3 67 Holmium Ho [Xe]4f10 5d1 6s2 +3 68 Erbium Er [Xe]4f115d1 6s2 +3 69 Thulium Tm [Xe]4f12 5d1 6s2 +3 70 Ytterbium Yb [Xe]4f14 5d0 6s2 +3, +2 71 Lutetium Lu [Xe]4f14 5d1 6s2 +3