Published on January 7, 2014
Periodic table and periodic properties Unit III - Classification of Elements and Periodicity in Properties Johann Dobereiner classified elements in group of three elements called triads. In Dobereiner’s triad the atomic weight of the middle element is very close to the arithmetic mean of the other two elements. Dobereiner’s relationship is referred as Law of triads. Since Dobereiner’s Law of triads worked only for few elements, it was dismissed. Chancourtois arranged elements in order of increasing atomic weights and made a cylindrical table of elements. John Newland arranged the elements in the increasing order of atomic weight and noted that the properties of the every eighth element are similar to the first one. This relationship is known as “Law of octaves”. Lothar Meyer proposed that on arranging the elements in order of increasing atomic weights similarities appear at a regular interval in physical and chemical properties. According to Mendeleev’s periodic law the physical and chemical properties of elements are periodic functions of their atomic weights. Merits of Mendeleev’s periodic table: Mendeleev’s periodic table was very helpful in remembering and studying the properties of large number of elements Mendeleev’s periodic table helped in correcting the atomic masses of some of the elements like gold, beryllium and platinum based on their positions in the periodic table. Mendeleev could predict the properties of some undiscovered elements like scandium, gallium and germanium. By this intuition, he had left gaps for the undiscovered elements while arranging elements in his periodic table. Demerits of Mendeleev’s periodic table: Position of hydrogen is not correctly defined in periodic table. It is placed in group I though it resembles both group 1 and 17. In certain pairs of elements increasing order of atomic masses was not obeyed. For example argon (Ar, atomic mass 39.9) is placed before potassium (K, atomic mass 39.1). Isotopes were not given separate places in the periodic table although Mendeleev's classification is based on the atomic masses. Some similar elements are separated and dissimilar elements are grouped together. For example copper and mercury resembled in their properties but had been placed in different groups. On the other hand lithium and copper were placed together although their properties are quite different. Mendeleev did not explain the cause of periodicity among the elements. Lanthanides and actinides were not given a separated position in the table. Modern Periodic Table: According to Modern periodic law the physical and chemical properties of the elements are periodic functions of their atomic numbers. Modern periodic table is also referred to as long form of periodic table. Horizontal rows in the periodic table are called periods. Vertical columns in the periodic table are called groups. In the modern periodic table there are 7 periods and 18 groups. The period number corresponds to highest principal quantum number of elements. First period contains 2 elements. – also called as shortest period Second and third period contains 8 elements. – called as shorter periods Fourth and fifth period contains 18 elements. – called as longer periods Sixth period contains 32 elements. – called as longest period In the modern periodic table, 14 elements of both sixth and seventh periods i.e. lanthanides and actinides respectively are placed separately at the bottom of the periodic table. Elements with atomic number greater than 92 are called transuranic elements. According to IUPAC, until a new element’s discovery is proved and its name is officially recognized it is given a temporary name. This nomenclature is based Latin words for their numbers.
Periodic table and periodic properties Classification of elements into blocks: The modern periodic table is divided into four main blocks – s -block, p-block, d-block and f-block depending on the type of orbital that are being filled with exception of hydrogen and helium. The elements in which last electron enter the s-orbital of their outermost energy level are called s-block elements. The s-block consists of two groups, Group-1 and Group-2. The elements of Group-1 are called alkali metals and have ns1 as the general outer electronic configuration. The elements of Group-2 are called alkaline earth metals and have ns2 as the general outer electronic configuration. The elements in which last electron enter the p-orbital of their outermost energy level are called p-block elements. The p-block elements constitute elements belonging to group 13 to 18. Elements of s-block and p-block are collectively called representative element The outermost electronic configuration of p-block elements varies from ns2np1 to ns2np6. Elements of group 18 having ns2np6 configuration are called noble gases. Elements of group 17 are called halogens. Elements of group 16 are called chalcogens. Number of valence electrons in group =Group number -10 for elements belonging to group 13 to 18. 36. Elements in which the last electron enters d-orbitals of penultimate energy level constitute d-block elements. Elements of group 3 to 12 in the centre of periodic table constitute the d-block elements. General outer electronic configuration of d-block elements is (n-1) d 1-10 ns 1-2. d-block elements constitute transition series elements. The name “transition series” is derived from the fact the d-block elements represent transition in character from reactive metals (belonging to group1 and 2 constituting s-block) on one side of the periodic table to non-metals (belonging to group 13 to 18 constituting p-block) on other side of the periodic table. Elements in which last electron enters f-orbitals are called f-block elements. Elements of Lanthanide series have general outer electronic configuration of 4f 1-14 5d 0-1 6s2. Elements of Actinide series have general outer electronic configuration of 5f 1-14 6d 0-1 7s2. Elements in lanthanide and actinide series are called inner transition series.
Periodic table and periodic properties Periodicity: The recurrence of similar properties of elements after certain regular intervals when they are arranged in order of increasing atomic number is called periodicity. The cause of periodicity of properties of elements is due to the repetition of similar electronic configuration of their atoms in the outermost energy shell after certain regular interval. Effective Nuclear Charge: • In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. • The nuclear charge that an electron experiences depends on both factors. • Effective nuclear charge is the net positive charge experienced by an electron. • The effective nuclear charge is not the same as the charge on the nucleus because of the effect of the inner electrons. • The electron is attracted to the nucleus, but repelled by the inner-shell electrons that shield or screen it from the full nuclear charge. • The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of electrons in the • • • spherical volume out to the electron in question. As the average number of screening electrons (S) increases, the effective nuclear charge (Zeff) decreases. Zeff = Z – S As the distance from the nucleus increases, S increases and Zeff decreases. S is called the screening constant which represents the portion of the nuclear charge that is screened from the valence electron by other electrons in the atom. The value of S is usually close to the number of core electrons in an atom. • Metal, Non-metals and Metalloids: • Metallic character refers to the extent to which the element exhibits the physical and chemical properties of metals. • Metallic character increases down a group. • Metallic character decreases from left to right across a period. Metals: • Metals are shiny and lustrous, malleable and ductile. • Metals are solids at room temperature (exception: mercury is liquid at room temperature; gallium and caesium melt just above room temperature) and have very high melting temperatures. • Metals tend to have low ionization energies and tend to form cations easily. • Metals tend to be oxidized when they react. • Compounds of metals with non-metals tend to be ionic substances. • Metal oxides form basic ionic solids. • Most metal oxides are basic: Metal oxide + water ------> metal hydroxide : Na2O(s) + H2O (l) ------> 2NaOH (aq) • Metal oxides are able to react with acids to form salts and water: Metal oxide + acid ------> salt + water: MgO(s) + 2HCl (aq) ------> MgCl2 (aq) + H2O (l) Non-metals: • Non-metals are more diverse in their behaviour than metals. • In general, non-metals are non-lustrous, are poor conductors of heat and electricity, and exhibit lower melting points than metals. • Seven non-metallic elements exist as diatomic molecules under ordinary conditions: • H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(l), I2(s) • When non-metals react with metals, non-metals tend to gain electrons: • • Metal + non-metal ------> salt 2Al(s) + 3Br2 (l) ------> 2AlBr3 (s) Compounds composed entirely of non-metals are molecular substances. Most non-metal oxides are acidic: Non-metal oxide + water ------> acid P4O10(s) + 6H2O (l) ------> 4H3PO4 (aq)
Periodic table and periodic properties CO2 (g) + H2O (l) ------> H2CO3 (aq) • Non-metal oxides react with bases to form salts and water: Non-metal oxide + base ------> salt + water CO2 (g) + 2NaOH (aq) ------> Na2CO3 (aq) + H2O (l) Metalloids: • Metalloids have properties that are intermediate between those of metals and non-metals. • Example: Si has a metallic lustre but it is brittle. • Metalloids have found fame in the semiconductor industry. • In general metallic character increases down the group and decreases along period. • In general non-metallic increases along a period and increases along group. Periodic Trends in Atomic Radii: • Atomic size varies consistently through the periodic table. • As we move down a group the atoms become larger. • As we move across a period atoms become smaller. • There are two factors at work: • The principal quantum number, n, and • The effective nuclear charge, Zeff. • As the principal quantum number increases (i.e., we move down a group), the distance of the outermost electron from the nucleus becomes larger. Hence the atomic radius increases. • As we move across the periodic table, the number of core electrons remains constant, however, the nuclear charge increases. Therefore, there is an increased attraction between the nucleus and the outermost electrons. This attraction causes the atomic radius to decrease. Periodic Trends in Ionic Radii: • Ionic size is important in predicting lattice energy. • In determining the way in which ions pack in a solid. • Just as atomic size is periodic, ionic size is also periodic. • In general, Cations are smaller than their parent atoms. • Electrons have been removed from the most spatially extended orbital. • The effective nuclear charge has increased. • Therefore, the cation is smaller than the parent atom. • Anions are larger than their parent atoms. • Electrons have been added to the most spatially extended orbital. • This means total electron-electron repulsion has increased. • Therefore, anions are larger than their parent atoms. • For ions with the same charge, ionic size increases down a group. • All the members of an isoelectronic series have the same number of electrons. • As nuclear charge increases in an isoelectronic series the ions become smaller: • Covalent radius for a homo nuclear molecule is defined as one half of the distance between the centres of nuclei of two similar atoms bonded by single covalent bond.
Periodic table and periodic properties • • • For hetero nuclear molecule covalent radius may be defined as the distance between the centre of nucleus of atom and mean position of the shared pair of electrons between the bonded atoms. Metallic radius is defined as the one half of the inter-nuclear distance two neighbouring atoms of a metal in a metallic lattice. For simplicity term atomic radius is used for both covalent and metallic radius depending on whether element is nonmetal or a metal. Comparison of atomic and ionic radii: Vanderwaal’s Radius: It is used for molecular substances in the solid state only. It is half of the distance between the nuclei of two adjacent nonbonded atoms in neighbouring molecules. Vanderwaal's radius is greater than the covalent radius as the Vanderwaal's forces are weak. Vanderwaal's radius is approximately 40% greater than the covalent radius Isoelectronic species: Atoms and ions which contain the same number of electrons, we call them isoelectronic species. E.g: O2–, F–, Na+ and Mg2+ have the same number of electrons (10). Their radii would be different because of their different nuclear charges. The cation with the greater positive charge will have a smaller radius because of the greater attraction of the electrons to the nucleus. Anion with the greater negative charge will have the larger radius, e.g. O 2–> F–> Na+> Mg2+ Lanthanide Contraction: In inner transition elements the differentiating electron enters into ‘f’ orbitals of the antepenultimate shell. As the atomic number increases in lanthanides due to the dispersed shape of f-orbitals and their poor shielding effect the atomic and ionic radii steadily decrease. This is called lanthanide contraction. Lanthanide contraction is also observed in 5d transition series. The atomic radius of 5d transition elements are very close to those of 4d transition elements due to Lanthanide contraction. As a result 4d and 5d transition elements are more similar in properties when compared to 3d and 4d transition elements E.g. Zr and Hf resemble most closed to each other than other elements.
Periodic table and periodic properties Ionization Energy: • The ionization energy of an atom or ion is the minimum energy required to remove an electron from the ground state of the • • • • isolated gaseous atom or ion. The first ionization energy, I1, is the amount of energy required to remove an electron from a gaseous atom: Na(g) ------> Na+ (g) + e– The second ionization energy, I2, is the energy required to remove the second electron from a gaseous ion: Na+ (g) ------> Na 2+ (g) + e– The larger the ionization energy, the more difficult it is to remove the electron. There is a sharp increase in ionization energy when a core electron is removed. Screening or shielding effect: The effective nuclear charge experienced by a valence electron in an atom will be less than the actual charge on the nucleus because of “shielding” or “screening” of the valence electron from the nucleus by the intervening core electrons. For example, the 2s electron in lithium is shielded from the nucleus by the inner core of 1s electrons. As a result, the valence electron experiences a net positive charge which is less than the actual charge of +3. In general, shielding is effective when the orbitals in the inner shells are completely filled. E.g: alkali metals which have a lone ns-outermost electron preceded by a noble gas electronic configuration Variations in Successive Ionization Energies: • Ionization energies for an element increase in magnitude as successive electrons are removed. • As each successive electron is removed, more energy is required to pull an electron away from an increasingly more positive ion. • A sharp increase in ionization energy occurs when an inner-shell electron is removed.
Periodic table and periodic properties • Ionization potential depends on the following factors. 1) Atomic radius. 2) Nuclear positive charge 3) Screening effect or shielding effect. 4) Extent of penetration of valence electrons. 5) Completely or half-filled sub shells. • • • With the successive removal of electrons ionization potential increases due to increased nuclear charge. I1 = 13.6 X z2 (z = effective nuclear charge) Ionization potential is measured in eV/atom and ionization energy is measured in kJ/mole (1 eV/atom = 96.45 kJ mole–1 or 1 eV/atom = 23.06 kcal/mole) • Ionization energies are determined from spectral studies as well as from discharge tube experiments. Periodic Trends in First Ionization Energies: • Ionization energy decreases down a group. • This means that the outermost electron is more readily removed as we go down a group. • As the atom gets bigger, it becomes easier to remove an electron form the most spatially extended orbital. • Example: For the noble gases the ionization energies follow the order He > Ne > Ar > Kr > Xe • The representative elements exhibit a larger range of values for I1 than transition metals. • Ionization energy generally increases across a period. • As we move across a period, Zeff increases, making it more difficult to remove an electron. • Two exceptions: removing the first p electron and removing the fourth p electron. • The s electrons are more effective at shielding than p electrons. So, forming the s2p0 configuration is more favourable. • When a second electron is placed in a p orbital, the electron-electron repulsion increases. When this electron is removed, the resulting s2p3 configuration is more stable than the starting s2p4 configuration. Therefore, there is a decrease in ionization energy. Electron Affinity: • • • • • • • • • Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion: Electron affinity and ionization energy measure the energy changes of opposite processes. Electron affinity: Cl(g) + e– ------> Cl– (g) ΔE = –349 kJ/mol Ionization energy: Cl(g) ------> Cl+ (g ) + e– ΔE = 1251 kJ/mol Electron affinity can either be exothermic (as the above example) or endothermic: Ar (g) + e– ------> Ar– (g) ΔE > 0 Look at electron configurations to determine whether electron affinity is positive or negative. The extra electron in Ar needs to be placed in the 4s orbital which is significantly higher in energy than the 3p orbital. The added electron in Cl is placed in the 3p orbital to form the stable 3p6 electron configuration. Electron affinities do not change greatly as we move down in a group Electron affinity: Electron affinities in for the s- and p-block elements in the first five rows of the periodic table. The more negative the electron affinity, the greater the attraction of the atom for an electron. An electron affinity > 0 indicates that the negative ion is higher in energy than the separated atom and electron.
Periodic table and periodic properties Electronegativity: It is defined as the ability of an atom or element in a molecule to attract the bond pair of electrons towards its direction. Mullikan in (1896-1986) defined electronegativity as proportional to the arithmetic mean of the ionization energy and the electron affinity. E.N= (I.E + E.A) / 2 Where E.N= Electronegativity I.E = Ionization Energy E.A = Electron Affinity Pauling Electronegativity: Linus Pauling was the original scientist to describe the phenomena of electronegativity. The best way to describe his method is to look at a hypothetical molecule that we will call XY. By comparing the measured X-Y bond energy with the theoretical X-Y bond energy (computed as the average of the X-X bond energy and the Y-Y bond energy), we can describe the relative affinities of these two atoms with respect to each other. Δ Bond Energies = (X-Y) measured – (X-Y) expected If the electronegativity of X and y are the same, then we would expect the measured bond energy to equal the theoretical (expected) bond energy and therefore the Δ bond energies would be zero. If the electronegativities of these atoms are not the same, we would see a polar molecule where one atom would start to pull electron density toward itself, causing it to become partially negative. By doing some careful experiments and calculations, Pauling came up with a slightly more sophisticated equation for the relative electronegativities of two atoms in a molecule: EN(X) - EN(Y) = 0.102 (Δ1/2). In that equation, the factor 0.102 is simply a conversion factor between kJ and eV to keep the units consistent with bond energies. By assigning a value of 4.0 to Fluorine (the most electronegative element), Pauling was able to set up relative values for all of the elements. This was when he first noticed the trend that the electronegativity of an atom was determined by its position on the periodic table and that the electronegativity tended to increase as you moved left to right and bottom to top along the table. The range of values for Pauling's scale of electronegativity ranges from Fluorine (most electronegative = 4.0) to Francium (least electronegative = 0.7). Furthermore, if the electronegativity difference between two atoms is very large, then the bond type tends to be more ionic, however if the difference in electronegativity is small then it is a nonpolar covalent bond.
Periodic table and periodic properties Allred-Rochow Electronegativity Allred and Rochow were two chemists who came up with the Allred-Rochow Electronegativity values by taking the electrostatic force exerted by effective nuclear charge, Zeff, on the valence electron. To do so, they came up with an equation: XAR = (3590 x Zeff/r2cov) + 0.744 At the time, the values for the covalent radius, rcov, were inaccurate. Allred and Rochow decided to add certain perimeters so that it would be more accurate and correspond to Pauling's electronegativity scale Trend of electronegativity: Electronegativity increases going from left to right across a period and decreases going down the group for the main group elements. Electronegativity is very useful in understanding the chemical properties of the elements. Valency: The combining capacity of the element in a molecule is called its valency • • • • The electrons present in the outermost shell are called valence electrons and these electrons determine the valence of atom. Valence of metallic element is usually equal to the group number of the group number of the element (or) the number of electrons in the valence shell. In the case of non – metals, it is equal to 8-(the number of electrons in the valence shell) On moving down a group since the number of valence electrons remains the same, all elements exhibit same valence. Oxidation number: Def: The apparent charge present on an atom of an element in a molecule is called oxidation number. For 1st group elements common oxidation number is +1 For 2nd group elements +2 For 3rd group elements +3 For 4th group elements +4 or -4 For 5th group elements -3 For 6th group elements -2 For 7th group elements -1
Periodic table and periodic properties Periodicity of the properties of oxides: Acidity and basicity of oxides of main group elements: From left to right in a period acidity of oxides increases From top to bottom of a group basicity of oxides increases. The oxides of group 13 elements are reactive to both an acid and a base and are named as amphoteric oxides. A well-known example is aluminium oxide Al2O3. Al2O3 + 6HCl ------> 2AlCl3 + 3H2O Al2O3 + 2NaOH + 3H2O ------>2Na [Al (OH) 4] Most of the oxides of non-metal elements are acidic. Their strength of acids increase going from left to right across a period of the periodic table. In other words, the acidity is stronger as the non-metal property increases. Periodicity in properties of hydrides: In going from left to right on the periodic table, the covalent character of hydrides decreases and the ionic character increases. Trend of Melting and Boiling Points: In a period MP's and BP's will increases first and decreases towards the end of period. In a group MP's and BP's will decreases from top to bottom in IA, IIA, IIIA, IVA. In VA, VIA, VIIA and zero groups, MP's and BP's will increases from top to bottom. TREND OF ATOMIC VOLUME: Atomic volume is the volume occupied by 1 gram – atom of element. In a group atomic volume increases from top to bottom. In a period atomic volume first decreases and then increases towards the end. TREND OF DENSITY: In a group density increases from top to bottom due to increase in atomic mass. In a period density increases from left to right due to decrease in atomic size
Periodic table and periodic properties Diagonal relationship: It is observed that some elements of second period show similarities with elements of third period present diagonally to each other though belonging to different group. This similarity in properties of elements present diagonally is called diagonal relationship. • Lithium is diagonally related to Magnesium, beryllium is diagonally related to aluminium and boron is diagonally related to silicon. • The anomalous behaviour of first element of s and p block elements of each group as compared to other group members is due to following reasons: • Small size of atom • Large charge/radius ratio • High electronegativity • Non availability of d-orbitals in their valence shell • In a periodic table there is high chemical reactivity at two extreme ends and lowest in centre. Maximum chemical reactivity at extreme left (alkali metals) is exhibited by the easy loss of electrons forming a cation and at extreme right (among halogens) shown by gain of electrons forming anion. • • The elements which readily loose electrons act as strong reducing agent. The element which readily accept electrons can act as strong oxidizing agent. The trend of Periodic Properties represented in a diagramatic manner:
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