CHEM 130 Chapter 3: Periodic Properties of Elements
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Module
CHEM 130 (CHEM130)
Institution
University Of Michigan - Ann Arbor
Class notes for Chapter 3: The Periodic Properties of Elements in the class General Chemistry: Macroscopic Investigations and Reaction Principles (CHEM 130) at the University of Michigan. Topics covered include the periodic table, electron spin and configuration, Coulomb's law, orbitals, ionization...
Aluminum: Low-Density Atoms Result in Low-
Density Metal
●Aluminum’s density is 2.70g/cm3
○In comparison iron is 7.86g/cm3, platinum is 212.4g/cm3
●Density of aluminum is low because the density of an aluminum atom is low ●Arrangements and mass-to-volume ratio are very important to consider when evaluating density
●Density of elements tends to increase as we move down a column in the periodic table
○Mass of each successive atom increases more than its volume does
●Periodic Properties: property of an element that is predictable based on an element’s position in the periodic table
●Structure determines properties The Periodic Table and Periodic Law
●Periodic Law: when elements are arranged in order of increasing mass, sets of properties recur periodically (found patterns)
●Modern periodic table organized by number of protons (atomic number)
○Most consistent way to see the trends of the properties
●Elemental properties are consistent down groups (columns)
○Main-group elements : elements whose properties are predictable based on position in the table
■Include groups 1A-8A
■If you number them 1-18, main group is 1, 2, 13-18
○Transition metals : properties tend to be less predictable based on position in periodic table
■Include groups 1B-8B, or 3-12
●Family/Group: one column within the main group of elements, all elements in a family exhibit similar chemical properties
Electron Configurations
●Electron Configuration: shows particular orbitals that electrons occupy for that atom
○EX: H → 1s1
■S → orbital, 1 → number of electrons in the orbital
●Ground State: lowest energy state of an atom or molecule
○Electrons generally occupy the lowest energy level available Electron Spin and the Pauli Exclusion Principle
●Orbital Diagram: diagram that symbolizes and electron as an arrow in a box representing an orbital, with arrow’s direction denoting electron’s spin
○
■Arrow is pointing up → m s = +½
●Pauli Exclusion Principle: no 2 electrons in the same atom can have the same 4 quantum numbers
○Each orbital can hold a maximum of 2 electrons and they will have opposing spins
○EX: helium has 2 electrons in the same orbital (1s) in the ground state
■
nLmLms
Electron 1100+½ Electron 2100-½
■
■Three quantum numbers in common bc same orbital, but different spin numbers
Sublevel Energy Splitting in Multi-Electron Atoms
●In a hydrogen atom the energy of an orbital only depends on n
○For hydrogen, the subshells within an energy level (n) are degenerate: they have
the same energy ■2s and 2p are degenerate, 3s 3p and 3d are degenerate, etc because H only has 1 electron – all other orbitals are empty
●In a multi-electron atom, they are not degenerate and energy depends on value of L
○Energies of sublevels are split
○Lower the value of L within a principal level, the lower the energy of the corresponding orbital
■For a given value of n → E ns < Enp < End < Enf
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