Chapter 8: Periodic Properties of the Elements Lizzie Rosenzweig

Periodic Property Periodic property is a property that is predictable based on an element’s position within the periodic table Ex. atomic radius, ionization energy and electron affinity The arrangement of elements in the periodic table reflects how electrons fill quantum mechanical orbitals

The Development of the Periodic Table Johann Döbereiner (1780-1849) German chemist Grouped elements into triads (threes with similar properties) John Newlands (1837-1898) English chemist Grouped elements into octaves (in analogy to music notes), these were groups of eight that had similar properties

The Development of the Periodic Table (cont) Dmitri Mendeleev (1834-1907) Russian chemist Periodic Law is when elements are arranged in order of increasing mass, and because of that, certain properties recur periodically The elements were arranged in increasing mass from left to right so, elements with the same properties fall into the same columns Mendeleev’s organization lead to the prediction of undiscovered elements and their properties There were some problems with atomic mass order Henry Moseley (1887-1915) Showed elements arranged in atomic number

Electron Configuration Electrons exist in orbitals Electron configuration shows the particular orbitals that electrons occupy for that atom Electrons generally occupy the lowest energy level orbitals available Ground state is the lowest energy state

Electron Spin and the Pauli Exclusion Principle An orbital diagram symbolizes the electron as an arrow and the orbital as the box The direction of the arrow represents the electron’s spin Electron’s spin is quantized as mₛ=½ or mₛ=-½ spinning up or down, respectively Pauli exclusion principle: no two electrons in an atom can have the same four quantum numbers This principle implies that each orbital can have a maximum of only two electrons, with opposing spins

Sublevel Energy Splitting in Multielectron Atoms When orbitals have the same energy they are called degenerate The energy of an orbital depends on the value of l The energies of the sublevels are split The lower the l value, the lower the energy of the corresponding orbital s<p<d<f

Coulomb’s Law Coulomb’s Law describes the interactions between charged particles Coulomb’s Law states the potential energy (E) of two charged particles depends on their charges (q₁ and q₂) and on their separation The amount of potential energy depends inversely on the separation between the charged particles

Conclusions from Coulomb’s Law For like charges, the potential energy is positive and decreases as the particles get farther apart. Since systems move toward lower potential energy, like charges repel each other. For opposite charges, the potential energy is negative and becomes and becomes more negative as the particles get closer together. Therefore, opposite charges attract each other. The interaction between charged particles increases in size as the charges of the particles increases

Shielding Shielding describes how one electron can shield another electron from the full charge of the nucleus Shielding is the repulsion of one electron by another electrons, screening that electron from the full effects of nuclear charge Effective nuclear charge is the actual nuclear charge experienced by an electron, defined as the nucleus plus the charge of the shielding electrons Inner electrons shield the outer electron from full nuclear charge

Penetration Penetration describes how one atomic orbital can overlap spatially with another, thus penetrating into a region that is close to the nucleus Penetration is when an outer electron occupies the region with the inner electrons and experiences full charge

Electron Spatial Distributions and Sublevel Splitting Splitting is a result of the spatial distributions of electrons within a sublevel Because of penetration, the sublevels of each principal level are not degenerate for multielectron atoms In the fourth and fifth principal levels, the effects of penetration become so important that the 4s orbital lies lower in energy than the 3d orbitals and the 5s orbital lies lower in energy that the 4d orbitals The energy separations between one set of orbitals and the next become smaller for 4s orbitals and beyond

Electron Configurations for Multielectron Atoms Aufbau principle is the pattern of orbital filling from lowest energy level to highest Hund’s rule states that when filling degenerate orbitals, a single electron fills the orbitals first Summarizing orbital filling Electrons occupy orbitals so as to minimize the energy of the atom; therefore, lower energy orbitals fill before higher energy orbitals One orbital has two electrons that spin in opposite directions When orbitals of identical energy are available, electrons first occupy these orbitals singly with parallel spins rather than pairs

Sublevels The s sublevel has 1 orbital and can hold 2 electrons The p sublevel has 3 orbitals and can hold 6 electrons The d sublevel has 5 orbitals and can hold 10 electrons The f sublevel has 7 orbitals and can hold 14 electrons

Practice Problems Write electron configurations for each element Mg P Br Al 1s² 2s² 2p⁶ 3s² 1s² 2s² 2p⁶ 3s² 3p³ 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵ 1s² 2s² 2p⁶ 3s² 3p¹

Practice Problem Write the orbital diagram for sulfur and determine the number of unpaired electrons. 1s² 2s² 2p⁶ 3s² 3p⁴ Two unpaired electrons