Periodic Properties of the Elements
See included chemistry topics below:
Effective Nuclear Charge, Atomic Size
periodic table development - ordered according to atomic mass by Mendeleev/Meyer
 Mendeleev
got majority of credit for advancing his ideas more vigorously and
predicting the existence of substances yet to be discovered
 Henry Moseley - arranged atoms by atomic number instead of atomic mass,
solved the problems in the original periodic table
effective nuclear charge - electric field created by nucleus and surrounding electron
density
 uses average environment created by nucleus/electrons
 Zeff = Z - S
o Z = number of protons in the nucleus
o S = average number of core electrons
 inner electrons shield outer electrons from the nucleus' charge
 charge increases as you move across any row/period of the periodic table
o Z increases as S stays the same, so Zeff increases
 charge increases only slightly as you move down a column/family
o larger electron cores less able to shield outer electrons than smaller
cores
atomic radii - nonbonding/bonding radius
 nonbonding radii (van der Waals radii) - radius of atoms not in molecules
 bonding radii (covalent radii) - radius of atoms in bonds
o slightly shorter than nonbonding radii
 increases as you move down column/family
o outer electrons get farther from nucleus as principal quantum number
increases
 decreases as you move left to right across a row/period
o as Zeff increases, outer electrons get pulled in more
 ionic radii - depends on charge
o cations - smaller than neutral atoms
o anions - larger than neutral atoms
 isoelectronic series - series of ions w/ same number of electrons,
number of protons
o radius decreases as nuclear charge increases
different
Ionization Energy, Electron Affinities
ionization energy - measures amount of energy needed to lose an electron
 more difficult to remove electrons w/ greater ionization energy
 more energy needed to remove each subsequent electron
 sharp increase in ionization energy needed to remove inner shell electrons
o inner shell electrons much closer to nucleus
 I1 generally increases w/ atomic number on each row
 I1 decreases as atomic number increases down a group
 representative elements have larger range of I1 than transition elements
 smaller atoms tend to have higher ionization energies (electrons closer to nucleus)
ion electron configurations - electrons removed from largest available quantum number
first
 electrons
added to lowest available quantum number first
electron affinity - energy change when electron is added to a gaseous atom
 measures attraction of atom for added electron
 usually negative (energy usually released when
electron is added), but can be
positive for noble gases (anion higher in energy than separated atom/electron)
 halogens have the most negative electron affinities
 noble gases have positive electron affinities (when adding an electron would place
it on a new energy subshell)
 group 5A (w/ 1/2 filled subshells) have electron affinities either positive or less
negative than group 4A
 doesn't change much going up/down a group
Metals, Nonmetals, Metalloids
metals - make up 3/4 of the periodic table, in left/middle part
 shiny luster
 malleable (can be pounded into thin sheets)/ductile (can be drawn into wires)
 good conductors of heat/electricity
 tend to form cations (have low ionization energies) in aqueous solutions
 compounds of metals w/ nonmetals tend to be ionic
 most metal oxides are basic (form hydroxides when dissolved in water)
 metal oxides w/ acid form salt
 alkali metals (group 1A) - soft metallic solids
o have the lowest I1 on each row
o hydride ion - H-; bonds w/ alkali metals to form hydrides
o
o
o
reacts exothermically w/ water
superoxide - O2-; combines w/ potassium, rubidium, cesium
emit certain colors when heated by a flame
 alkaline earth metals (group 2A) - denser/harder than alkali metals
o less reactive than alkali metals
o Mg used for lightweight structure alloys because layer of MgO protects
it from other chemicals
nonmetals - on right side of periodic table
 no luster, has various appearances
 poor conductors of heat/electricity
 most nonmetal oxides are acids
 tend to form anions/oxyanions in aqueous solutions
 molecular substances - compounds made up only of nonmetals
 hydrogen - doesn't really belong in a particular group
o exists as H2 gas in most conditions
o can be metallic at very high pressures
o has much higher I1 than other alkali metals (lacks any type of
nuclear
shielding)
 oxygen group (group 6A)
o elements change from nonmetal to metal as you go down the group
o allotrope - different form of same element in same state
o ozone - O3, less stable than O2
o sulfur - exists naturally as S8 rings
 halogens (group 7A) - aka "salt formers"
o melting/boiling points increase as atomic number increases
o very highly negative electron affinities
o Cl - most industrially useful halogen
o forms halide compounds w/ hydrogen
 noble gases (group 8A) - monoatomic nonmetals, gas at room temperature
o have completely filled s and p subshells
o very unreactive
metalloids - have properties intermediate between metals/nonmetals
 have only some metallic properties, but lack others
 many used as electrical semiconductors, integrated circuits