University:
Bradley University
Course:
CHM 100 | Fundamentals of General Chemistry
Academic year:

2020
Views:

109

Pages:

9
Author:

Gudrunecg9m

IV. Electron Configurations: Valence Electrons and the Periodic Table We saw that Mendeleev arranged elements with similar periodic properties in the same column. Also, notice that as you move down a column, the number of electrons in the outermost principle energy level stays the same. The key connection between the macroscopic world (an element’s chemical properties) and the microscopic world (an atom’s electronic structure) lies in the outermost electrons. • VALENCE ELECTRONS: - Located in highest energy s & p orbitals - The e− involved in bonding / ion formation - Determine element’s chemical properties FOR ELEMENTS IN GROUPS 1 – 8: # valence e─ = group # FOR TRANSITION METALS: count the highest energy s, p, and d e─ • CORE ELECTRONS: inner e─ ; exist in completely filled energy levels IV. The Explanatory Power of the Quantum- Mechanical Model Electron configurations from the periodic table: 1 1s1 2 2s1 2s2 3 3s1 3s2 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 7d 1s2 4f 5f 6f 7f 2p1 2p2 2p3 2p4 2p5 2p6 3p1 3p2 3p3 3p4 3p5 3p6 4 4s1 4s2 3d1 3d2 3d3 3d4 3d5 3d6 3d7 3d8 3d9 3d10 4p1 4p2 4p3 4p4 4p5 4p6 4 5 6 5 5s1 5s2 4d1 4d2 4d3 4d4 4d5 4d6 4d7 4d8 4d9 4d10 5p1 5p2 5p3 5p 5p 5p 6 7 6s1 6s2 5d1 5d2 5d3 5d4 5d5 5d6 5d7 5d8 5d9 5d10 6p1 6p2 6p3 6p4 6p5 6p6 7s1 7s2 6d1 6d2 6d3 6d4 6d5 6d6 6d7 6d8 6d9 6d10 7p1 7p2 7p3 7p4 7p5 7p6 s d p 4f1 4f2 4f3 4f4 4f5 4f6 4f7 4f8 4f9 4f10 4f11 4f12 4f13 4f14 5f1 5f2 5f3 5f4 5f5 5f6 5f7 5f8 5f9 5f10 5f11 5f12 5f13 5f14 f *Because elements in the same column of the periodic table have the same number of valence electrons, they have similar properties (reactivity, ionic charge). VI. Periodic Trends in the Size of Atoms and Effective Nuclear Charge • ATOMIC RADIUS: - used to represent the size of an atom, assuming it is roughly spherical - measured as half the distance between two nuclei of bonded atoms TREND: DOWN a column: radii INCREASE RIGHT across a row: radii DECREASE EXPLAINING THE TREND:  Down a column: e─ occupy larger orbitals further from nucleus so e─ are more spread out  radius increases  Right across a row: e─ enter orbitals in the same energy level (equal distance from nucleus) ; nuclear charge also increases (an attractive force) so e─ are being pulled closer to the nucleus  radius decreases VII. Ions: Electron Configurations, Magnetic Properties, Ionic Radii, and Ionization Energy MAGNETIC PROPERTIES • PARAMAGNETIC: an atom/ion that contains unpaired e─ and is attracted by a magnetic field • DIAMAGNETIC: an atom/ion that contains no unpaired e─ and is not attracted by a magnetic field Note that an element may be diamagnetic but an ion of that element may be paramagnetic. In order to determine if an atom/ion is paramagnetic or diamagnetic, we must look at its orbital diagram. IONIC RADII • CATIONS: positive ions are smaller than their corresponding atoms  Ca2+ smaller than Ca WHY? • ANIONS: negative ions are larger than their corresponding atoms  F─ larger than F WHY? • ISOELECTRONIC: atoms/ ions w/ the same number of e─ Ex. O2─ and Na+ both have 10 e ─ To avoid confusion when comparing the size of + and ─ ions to each other: • Remember that + ions are always smaller b/c they become isoelectronic w/ elements on the far right side of the periodic table • Consider which elements the + ions are isoelectronic with and compare the location of these elements on the periodic table. IONIZATION ENERGY • IONIZATION ENERGY: energy required to remove an electron from an atom or an ion in its gaseous state Ex. Na (g)  Na+ (g) + 1 e─ IE1 = 496 kJ/mol TREND (1ST IONIZATION ENERGY): DOWN a column: IE generally DECREASES RIGHT across a row: IE generally INCREASES