Periodic Trends PDF

Course Outcome 6 CHEM 1102 (Principles of Chemistry Lab) Periodic Trends PAUL JHON G. EUGENIO Faculty, Department of Chemistry 2  Periodic Table ▸ List of all known elements ▸ Seven (7) rows and eighteen (18) columns ▸ Arrangemen...

Course Outcome 6 CHEM 1102 (Principles of Chemistry Lab) Periodic Trends PAUL JHON G. EUGENIO Faculty, Department of Chemistry 2  Periodic Table ▸ List of all known elements ▸ Seven (7) rows and eighteen (18) columns ▸ Arrangement shows regular pattern in terms of atomic properties 3  Historical Development ▸ Johann Wolfgang Döbereiner ▸ John Alexander Reina Newlands ▸ Dmitri Ivanovich Mendeleev ▸ Julius Lothar Meyer ▸ Henry Gwyn Jeffreys Moseley 4  Law of Triads J.W. Dobereiner ▸ The properties of the middle member of the triad is approximately equal to the average of the properties of the first and the third member. ▸ Examples of triads: Li, Na, K 5  Law of Octaves J.A.R. Newlands ▸ Eight musical notes ▸ The first has the same properties with the eight member ▸ The chemical elements are arranged according to increasing atomic weight, those with similar physical and chemical properties occur after each interval of seven elements. 6  Periodic Law ▸ When atoms are arranged in order of increasing atomic masses, atomic properties of the elements show a regular pattern. ▸ Proposed by Mendeleev and Meyer 7  Dmitri Mendeleev ▸ Father of Modern Periodic Table ▸ Arrange the elements in order of increasing atomic mass 8  Lothar Meyer ▸ Came up with the same scheme with that of Mendeleev ▸ Used molar volumes 9  Why Mendeleev was given the credit for the first modern periodic table? 10  Dmitri Mendeleev ▸ First to published his work ▸ Able to predict the existence of yet to be discovered elements ▸ Leaves spaces/blanks in his table ▸ Able to correct some of the computed atomic masses 11  Problem with Mendeleev’s Periodic Table ▸ Two different elements occupying the same cell/square in the table ▸ Ar and Ca ▸ Ar and K 12  Henry Moseley ▸ Revised the periodic law: ▸ When elements are arrange in order of increasing atomic numbers, their atomic properties show a regular interval or pattern. 13  Periodic Table ▸ Vertical Columns - represents groups or families ▸ Horizontal Rows - represents the period of the element Groups in the Periodic Table ▸ A – representative elements ▸ B - transition metals 14  Representative Elements 1A = Alkali metals 2A = Alkaline Earth Metals 3A = Boron Family 4A = Carbon Group 5A = Nitrogen Group / Pnictogens 6A = Oxygen Group / Chalcogens 7A = Halogens 8A = Noble Gas 15  Blocks of the Periodic Table ▸ s-block = group 1A and 2A ▸ p-block = from boron group to noble gas (groups 3A to 8A) ▸ d-block = transition metals ▸ f- block = inner-transition metals 16  Electron Configuration and Periodic Table ▸ Members of the group in the periodic table possess similar properties because of their similarity in terms of their valence electrons which is determine by electron configuration 17 18  Alkali Metals ▸ Li = 1s2 2s1 ▸ Na = 1s2 2s2 2p6 3s1 ▸ K = 1s2 2s2 2p6 3s2 3p6 4s1 19  Motivating Exercises 1 Identify the element having: 1. 1s2 2s2 2p6 Ne 2. 1s2 2s2 2p6 3s2 3p5 Cl 3. 1s2 2s2 2p2 C 4. [Ar] 4s2 3d5 Mn 5. [Kr] 5s2 4d10 5p3 Sb 20  Motivating Exercises 2 Locating an Element: Group and Period 1.Group 1A, period 4 K 2. Group 5A, period 3 P 3. Group 3A, period 5 ln 4. Group 2A, period 6 Ba 5. Group 8A, period 2 Ne 21  Periodic Trends 1. Atomic /Ionic Size 2. Ionization Energy 3. Electron Affinity 4. Electronegativity 22  Atomic Radius The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei. Trend: Top to Bottom: increases Left to Right: decreases 23  Atomic Radius What is the reason for such trend? Top to Bottom: energy level Left to Right: protons 24  Top to Bottom: energy level ▸ As you move down a group, energy level increases thus making the size larger ▸ Imagine an onion cut cross-sectional, the rings are the energy level 25  Left to right: protons ▸ As you move from left to right in a period, more protons inside the nucleus, the greater the force of attraction for its electrons, making closer to the nucleus. 26  Motivating Exercise 3 Arrange the following in order of increasing atomic radius: 1. Rb, Li, Na, K and Cs Li, Na, K, Rb, Cs 2. C, B, F, N and O F, O, N, C, B 27  Ionic Size Ionic size depends upon: 1. Nuclear charge. 2. Number of electrons. 3. Orbitals in which electrons reside. 28  Ionic Size Cations are smaller than their parent atoms. The outermost electron is removed and repulsions are reduced. 29  Ionic Size Anions are larger than their parent atoms. Electrons are added and repulsions are increased. 30  Ionic Size Rule ▸ The more positive the charge of the ion is, the smaller its size. ▸ The more negative the charge of the ion is, the larger its size. 31  Ionic Size In an isoelectronic series, ions have the same number of electrons. Ionic size decreases with an increasing nuclear charge. 32  Motivating Exercise 4 Choose the largest among the group +1 A. Na+1 , Al3+ , Mg2+ Na B. S-2, Cl-1 , P-3 -3 P 33  Motivating Exercise 5 Arrange the following in order of increasing ionic size. N3- O2- F-1 Ne Na+ Mg2+ 2+ + -1 2- 3- Mg Na Ne F O N 34  Ionization Energy ▸ Amount of energy required to remove an electron from the ground state of a gaseous atom or ion. ▸ First ionization energy is that energy required to remove first electron. ▸ Second ionization energy is that energy required to remove second electron, etc. Trend: Top to Bottom: decreases 35 Left to Right: increases  Ionization Energy ▸ It requires more energy to remove each successive electron. ▸ When all valence electrons have been removed, the ionization energy takes a quantum leap. 36  Ionization Energy ▸ going down a column, less energy to remove the first electron. ▹ For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus. 37  Ionization Energy ▸ Generally, it gets harder to remove an electron going across. ▸ As you go from left to right, Zeff increases. 38  Ionization Energy ▸ On a smaller scale, there are two jags in each line. Why? 39  Ionization Energy ▸ The first occurs between Groups IIA and IIIA. ▸ Electron removed from p-orbital rather than s-orbital ▹ Electron farther from nucleus ▹ Small amount of repulsion by s electrons. 40  Ionization Energy ▸ The second occurs between Groups VA and VIA. ▹ Electron removed comes from doubly occupied orbital. ▹ Repulsion from other electron in orbital helps in its removal. 41  Electron Affinity ▸ The negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion ▸ Energy needed for an atom to attract an electron Trend: Top to Bottom: decreases Left to Right: increases 42  Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row. The more negative the value for electron affinity corresponds to greater attraction for an electron 43  Electron Affinity There are also two discontinuities in this trend. 44  Electron Affinity The first occurs between Groups IA and IIA. ▹ Added electron must go in p-orbital, not s-orbital. ▹ Electron is farther from nucleus and feels repulsion from s-electrons. 45  Electron Affinity The second occurs between Groups IVA and VA. ▹ Group VA has no empty orbitals. ▹ Extra electron must go into occupied orbital, creating repulsion. 46  Electronegativity ▸ The ability of an atom to attract electrons toward itself in a chemical bond Trend: Top to Bottom: decreases Left to Right: increases 47