The Periodic Table and Trends
Topics 2 and 3
Please have a periodic table out.
SONG
IB prefers this one.
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
numbers (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
Know the names of these groups
• 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
2
1
3
2
5
8
What two blocks will always be the highest occupied level?
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 in
the ratio 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
This is a simple, yet very good picture. Do you understand it?
• ATOMIC RADII
– 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
– 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
McGraw » increases the strength of the positive
Hill
nucleus and pulls electrons in the given
video
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
IONIC RADII
Looking at ions compared to their
parent atoms
• atoms tend to gain or loose electrons in
order to have the electron configuration
of a noble gas
– do atoms become smaller or larger when
they do this?
– 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 forming a
cation
Li+
.078nm
– anions (- ions) are larger than parent atom
• have gained an electron 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 nm
9e- and 9p+
F- 0.133 nm
10 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
DARK GREY IS THE SIZE OF THE ION
– 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
• the increase in ionization energy is
not continuous across the table
• electrons are also harder to
remove…
–a sub-level (s,p,d,f) is completely filled
–a sub-level (s,p,d,f) are half filled
notice discontinuity as
move across period 3
only last sublevel shown
• ELECTRON AFFINITY (Ea)
– the change in energy (kJ/mole) of a
neutral atom (in the gaseous phase) when
an electron is added to the atom to form a
negative ion (anion)
– in other words, the neutral atom's
likelihood of gaining an electron
– example
• F(g) + e-  F-(g) will release 328 kJ/mole
of energy
– the more negative Ea, the greater the attraction
for the electron
• trends
– across a period
• in general, Ea become more negative from
L to R
– down a group
• in general, becomes less negative
This one same as “IB textbook”
• 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 electrons
H
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 a shared pair of electrons
as they will be in a level closer to the
nucleus moving from L to R on the table
• next concepts require understanding of
concepts covered in later topics (this year
and even senior year)
• only need to know the trends, not the
reason why until later
– MELTING POINT
• down group 1 (alkali metals)
– decreases as “sea of negative
electrons” are farther away
from the positive metal ions
• down group 17 (halogens)
Element
Melting
Point (K)
Li
453
Na
370
K
336
Rb
312
Cs
301
Fr
295
– increases as the van der Waals’ forces increase
» larger molecules have more
electrons which increases
the chance that one side of
the molecule could be negative
increases
increases
• across the table (period 3)
– from left to right
• increases until group 14 (think diamonds)
then decreases starting at group 15
– bonding goes from strong metallic to very
strong macromolecules (network covalent) to
weak van der Waals’ attraction
• CHEMICAL PROPERTIES
– groups
1+ charge
• alkali metals
– react vigorously with water and air
» 2Na (s) + H2O (l)  2Na (aq) + 2OH- (aq) + H2 (g)
» (Li, Na, K… all the same equation)
» reactivity increases downwards
» because the outer (valence) electron is in higher energy
levels (farther from the nucleus) and easier to remove
– react with the halogens
» halogens’ reactivity increases upwards
» smaller size attracts electrons better
since they can be close to the
nucleus
1- charge
least reactive
most reactive
• halogens (group 7)
–diatomic molecules such as Cl2, Br2, I2
» can react with halide ions (Cl -, Br -, and I -)
» the most reactive ends up as an ion (1- charge)
and is not visible (molecules Cl2, Br2, I2 are a
visible gas)
» Cl > Br > I
Cl-(aq)
Cl2
Br-(aq)
Colorless- no turns
red
due
reaction
formation of Br2
I-(aq)
to turns brown due
formation of I2
to
to
Br2 no reaction
no reaction
turns brown due
formation of I2
I2 no reaction
no reaction
no reaction
– periods
• from left to right in period 3
– metals…metaloids…nonmetals
– oxides are
» ionic…..and then covalent bonds
– when oxides react with water
» basic…amphoteric (either basic or
acidic)…acidic
»
»
»
»
Na2O(s) + H2O (l)  2 NaOH (aq) strong base
MgO (s) +H2O (l)  Mg(OH)2 (aq) weaker base
P4O10 (s) + 6H2O (l)  4 H3PO4 (aq) weak/strong acid
SO3(g) + H2O (l)  H2SO4 (aq) strong acid
Look at the blue arrows! Senior year…