Periodic Trends
Periodic Table Development
 1869- Mendeleev and Meyer published classification schemes for elements and observed that similar
chemical and physical properties occur periodically when elements are in order by atomic mass
 Mendeleev organized them and left blank spaces for undiscovered elements where he predicted their
properties
 Periodic Law = when elements are arranged in order of increasing atomic number, there is a periodic
repeating pattern to their properties
 Trends in properties are related to the electron configurations of elements and two phenomena:
shielding effect and nuclear charge!
http://video.google.com/videoplay?docid=-2134266654801392897
Shielding Effect and Nuclear Charge – explain periodic trends
 Shielding effect = inner electrons block valence (outer) electrons from the pull of the positively charged
nucleus
o The more principal energy levels, the more “layers of” inner electrons available to shield the
valence electrons
o From left to right on the table, shielding is constant (electrons are being added to the same
principal energy level- no new levels are added)
o From top to bottom on the table, shielding increases (more principal energy levels are being
added)
 Nuclear charge = pull or attraction for electrons by the nucleus
o The higher the atomic number, the more protons in the nucleus, and the stronger the pull of the
nucleus (greater nuclear charge)
o From left to right on the table, nuclear charge increases (more protons means more nuclear
charge)
All other periodic property trends result from these!
Periodic Trends (focusing on representative elements, Groups 1, 2, 13 -18)
Atomic radius (size of the atom)
 From left to right on the table, atomic radius DECREASES
o Why? Shielding is constant but nuclear charge increases
 From top to bottom on the table, atomic radius INCREASES
o Why? Shielding noticeably increases
Ionic radius (size of the ion)
 Cations are smaller than their parent atoms (fewer electrons compared to protons, electrons pulled in tighter
to nucleus)
 Anions are larger than their parent atoms (more electrons compared to protons, electrons farther out from
nucleus)
 Cation size decreases from left to right, anion size decreases from left to right (same reasons as atomic
radius)
 Cation size increases from top to bottom, anion size increases from top to bottom (same reasons as atomic
radius)
Ionization energy (energy required to remove an electron… first IE is the energy required to remove the
outermost valence electron, second IE is the energy to remove the next one, etc. )
 From left to right on the table, ionization energy INCREASES
o Why? Shielding is constant but nuclear charge increases
o If the electrons are being held tighter, it is harder to remove them (takes more energy)
 From top to bottom on the table, ionization energy DECREASES
o Why? Shielding noticeably increases
o If the electrons are not being held tightly, it takes less energy to remove them
 THINK about how the ionization energy might change depending on the electron configuration of the
atoms/ions we are examining (is first IE higher or is second, third, etc.)
Electronegativity (ability of an atom to attract other electrons to form a bond)
 From left to right on the table, electronegativity INCREASES
o Why? Shielding is constant but nuclear charge increases
o If the electrons of the atom are being pulled more strongly, outside electrons will also be pulled more
strongly
 From top to bottom on the table, electronegativity DECREASES
o Why? Shielding noticeably increases
o If the electrons of the atom are not being pulled strongly, outside electrons will not be either
Reactivity and Metallic Character
 Metals have a tendency to lose electrons when they react
 Nonmetals have a tendency to gain electrons when they react
 Elements with low ionization energies or high electronegativities will be the most reactive since they have
the strongest tendencies to lose and gain electrons, respectively (the more reactive elements are in the lower
left corner or upper right corner)
 From left to right on the table, metallic character DECREASES (electrons are not easily removed,
nonmetals are on the right side)
 From top to bottom on the left side of the table, metallic character INCREASES (electrons are more easily
lost)
Periodic Table Review
 Metals, nonmetals, metalloids
o Properties of metals
 Shiny luster
 Malleable and ductile
 Good conductors of heat and electricity
 Form cations (lose electrons)
 High melting points
o Properties of nonmetals
 Dull
 Brittle, hard, or soft
 Poor conductors of heat and electricity
 Form anions (gain electrons)
 Low melting points (some are liquids or gases at room temp)
o Properties of metalloids
 Have properties of both
 Can form cations or anions
 Are often electrical semiconductors
 States of matter of elements (how can you tell?)
 Alkali metals, alkaline earth metals, transition metals or elements, halogens, noble or inert gases, inner
transition metals, rare earth elements, lanthanide series, actinide series (be able to find them!)