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The Periodic Table and Trends
SONGPlease have a periodic table out.
Dmitri Mendeleev 1834 – 1907
• Russian chemist and teacher• given the elements he knew
about, he organized a “Periodic Table” based on increasing atomic mass (it’s now atomic #)
• he even left empty spaces to be filled in later
At the time the elements gallium and germanium were not known. These are the blank spaces in his periodic
table. He predicted their discovery and estimated their properties.
Henry Moseley 1887 – 1915
• arranged the elements in increasing atomic number (Z)– properties now recurred
periodically
Design of the Table• Groups are the vertical columns.
– elements have similar, but not identical, properties• most important property is that
they have the same # of valence electrons
• valence electrons- electrons in the highest occupied energy level
• all elements have 1,2,3,4,5,6,7, or 8 valence electrons
These are a lower level. Therefore the d sub-level
is never included for valence electrons
The highest level is 4.
Lewis Dot-Diagrams/Structures
• a short cut wherevalence electrons are represented as dots around the chemical symbol for the element
Na
Cl
Look, they are following
my rule!
• B is 1s2 2s2 2p1;– 2 is the outermost energy level – it contains 3 valence electrons, 2 in
the 2s and 1 in the 2p• Br is [Ar] 4s2 3d10 4p5
How many valence electrons are present?
• Periods are the horizontal rows– do NOT have similar properties– however, there is a pattern to their properties
as you move across the table that is visible when they react with other elements
Trends in the table IB loves the alkali metals and the
halogens
• many trends are easier to understand if you comprehend the following
• the ability of an atom to “hang on to” or attract its valence electrons is the result of two opposing forces – the attraction between the electron and the
nucleus– the repulsions between the electron in
question and all the other electrons in the atom (often referred to the shielding effect)
– the net resulting force of these two is referred to effective nuclear charge
– the distance from the nucleus to the outermost electron
– cannot measure the same way as a simple circle due to electrons are not in a fixed location
– therefore measure distance between two nuclei and divide by two
• ATOMIC RADII
–groups • increases downwards as more levels are added• more shielding
–periods across the periodic table• radii decreases
–the number of protons in the nucleus increases» increases the strength of the positive
nucleus and pulls electrons in the given level closer to it
»added electrons are not contributing to the shielding effect because they are still in the same level
H
Li
Na
K
Rb
McGraw Hill video
– cations (+ ions) are smaller than the parent atom• have lost an electron (actually, lost an entire level!)• therefore have fewer electrons than protons
Li 0.152 nm
Li+ .078nm
+Li forming a
cation
–anions (- ions) are larger than parent atom• have gained an electron(s) to achieve noble gas
configuration• effective nuclear charge has decreased since
same nucleus now holding on to more electrons• plus, the added electron repels the existing
electrons farther apart (kind of “puffs it out”)
F 0.064 nm9e- and 9p+
F- 0.133 nm10 e- and 9 p+
-
– trends• across a period
–decreases at first when losing electrons (+ ion)–then suddenly increases when gaining electrons
(- ion)–then goes back to decreasing after just like
neutral atoms because of more protons pulling in the outer level
• down a group (same as neutral atoms)– increases as new levels are added–more levels shielding
Looking at ions compared to their parent atoms
• meaning does an atom become smaller or larger as it gains or loses electrons?
IONIC RADII
– IONIZATION ENERGY• the minimum energy (kJ mol-1) needed to
remove an electron from a neutral gaseous atom in its ground state, leaving behind a gaseous ion– X(g) X+(g) + e-
• first ionization energy- energy to remove first electron
• second ionization energy- energy to remove second electron
• third ionization energy- and so on…
don’t forget-- gaseous
• decreases down a group–outer electrons are farther from the nucleus and
therefore easier to remove– inner core electrons “shield” the valence electrons
from the pull of the positive nucleus and therefore easier to remove
• increases across a period– the nucleus is becoming stronger (effective
nuclear charge) and therefore the valence electrons are pulled closer• atomic radii is decreasing • this makes it harder to remove a valence electron
since it is closer to the nucleus
– or another way to look at it… a stronger nuclear charge acting on more contracted orbitals
• ELECTRONEGATIVITY–measures the attraction for a shared pair
of electrons in a bond• Linus Pauling (1901 to 1994) came up with a
scale where a value of 4.0 is arbitrarily
given to the most electronegative element,
fluorine, and the other electronegativities
are scaled relative to this value.
• trends (same as ionization energy and for the same reasons)
• as you go down a group electronegativity decreases – the size of the atom increases
» the bonding pair of electrons (-) is increasingly distant from the attraction of the nucleus (+)
» the bonding pair of electrons (-) are shielded because of core electrons (-) interfering with the nucleus’ (+) hold on valence electronsH
Li
Na
K
Rb
• as you go across a period– electronegativity increases
• the atoms become smaller as the effective nuclear charge increases –easier to attract the shared pair of
electrons as they will be in an orbital closer to the nucleus moving from L to R on the table
–MELTING POINT• down group 1 (alkali metals)
–decreases
• down group 17 (halogens)–decreases
Element Melting Point (K)
Li 453
Na 370
K 336
Rb 312
Cs 301
Fr 295
• across the table (period 3)– from left to right
• increases until group 14 (think diamonds) then decreases starting at group 15 (gases)