Chem Chapter 8 Electron Configurations and Periodicity PDF

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Chem 107 Binghamton University Chapter 8 Electron Configurations and Periodicity Aufbau Principle-When  electrons are in the lowest energy orbitals they’re the most stable -When electrons are added to an atom they go to the lowest energy orbital available -Each orbital can hold two electrons -Subshells fill in the order of spdf but they can overlap with different values of n, such as 4s>3d Degenerate-Orbitals that have the same energy level Core Electrons-Filled  inner shells in an atom that are not part of the chemical reactions Valence Electrons-in  the outermost shell and influence the atom’s chemical behavior the most Hund’s Rule-Degenerate  orbitals like 2p have the lowest configuration that maximizes the number of unpaired electrons → Filled shells can also be shown through using the noble gas [He]=1s^2 Exceptions Cr 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^5 [Ar]4s^1 3d^5 Cu 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^10 [Ar]4s^1 3d^10 -exceptions rise from half filled and completely filled d-subshells

Electron Configurations:Ions -Ions are the lose or gain of valence electrons to achieve a stable electron configuration Cations -Na(g) → Na+(g)+e-[Ne]3s^1 → [Ne]+e-Lose an electron and becomes positive Anions -Cl(g)+e- → Cl-(g) -[Ne]3s^2 3p^5 +e- → [Ne]3s^2 3p^6=[Ar] -Gains an electron Isoelectronic Atoms/Ions -Other main group elements are formed through gaining or losing electrons to create noble gases -Mg → Mg2+ +2e- =[Ne] -O+2e- → O2- = [Ne] -Atoms or ions that have the same electron configuration are isoelectronic -Ex. Na+, Mg2+, O2-, F-, Ne = [Ne] Cations of Transition Metals -The loss of valence electrons (s) or (d) to achieve stable configurations -Fe → [Ar]4s^2 3d^6 -Fe2+ → [Ar] 3d^6 (Loss the valence electrons)

-Fe3+ → [Ar]3d^5 (lost one 3d and is now a half filled subshell) *When 3d is full or partially full it is below the 4s level Magnetic Properties of Atoms -Only atoms with unpaired electrons show magnetism Paramagnetic Substance-one  that is weakly attracted to a magnetic field because they usually have unpaired electrons Diamagnetic Substance-not  attracted to a magnetic field because they only have paired electrons Sizes of Atoms/Ions Atomic Radius-the  distance between nuclei when bonded in molecules -Diatomic molecules: The covalent radius which is ½ the distance between nuclei -Metals:The metallic radius ½ the distance between nuclei in metal lattice -Ions:The ionic radius is ½ the distance between ions in ionic crystal lattice Atomic Radii -Increases going down a family because energy levels are being added -The electrons on the inside decrease the effect of nuclear charge -The atomic size increases as n increases -Decreases going across a row -The nuclear charge increases as it goes across a row -There is increased attraction to the electrons which causes the atomic size to decrease Effective Nuclear Charge -The positive charge an electron experiences from the nucleus which is equal to the nuclear charge but can be reduced by intervening electron distribution -Increases across a period because n is the same across and each has a stronger nuclear charge -The radius of the atom decreases with an increase of atomic number Radius of Ions -Cations: Lose electrons from their valence shell orbitals which decreases their radius -Anions: Gain electrons in their valence shell which increases their radius *Anions will always be larger than cations they’re isoelectronic with *If there is an element that is isoelectric but has more protons it has a smaller radius Ionization Energy -Is the quantity of energy needed to remove 1 mole of e- from 1 mole an atom or ion in the gaseous state (kj/mole) -X(g) → X+(g)+e-(g) -1st ionization energy: Mg(g) → Mg+(g)+e-2nd ionization energy: Mg+(g) → Mg2+(g)+e-

-The ionization energy increases from IE2>IE1 -Decreases as it goes down a family because it takes less energy to remove electrons from higher energy orbitals since the electrons are further away -Increases going across a row because it gets harder to remove electrons as the nuclear charge increases Electron Affinities -Energy released as 1 mole of gaseous atoms gain electrons to form anions -EA becomes more negative moving to the right and up in the periodic table...