Download THE PERIODIC TABLE

The Periodic Table and
Physical Properties
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Topics 3.1 - 3.3 and 12.1.1 - 12.1.2
Dmitri Mendeleev
8 February 1834 – 2 February 1907
• Russian chemist and teacher
• given the elements he knew
about, he organized a
“Periodic Table” based on
increasing atomic mass (it’s
atomic # now)
• he even left empty spaces to
be filled in later (TOK– he
was a “scientist” and “risk
taker”!)
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.
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
Electron arrangement (SL level – 3.1.3)
2
2,1
2,3
2,5
2,8
2,8,2
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• B is 1s2 2s2 2p1;
– 2 is the outermost energy level
– it contains 3 valence electrons, 2 in
the s and 1 in the p
• 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
Definitions
• atomic radii
– the distance from the nucleus to the outermost
electron
• ionic radii
– same distance, but for ions (atoms that have lost
or gained valence electrons)
• first ionization energy (kJ mol-1)
– the energy needed to remove the outermost, or
highest energy, electron from a neutral atom in
the gaseous phase
• electronegativity
– measures the attraction for a shared pair of
electrons
• melting point
• chemical properties
– how elements react with other elements
Trends in the table
But first, the electron shielding effect
• electrons
between the
nucleus and
the valence
electrons
repel each
other
• ATOMIC RADII
– McGraw Hill video
– groups (alkali metals and halogens)
H
Li
Na
K
• increases downwards as more levels are added
– periods across the periodic table (period 3)
• radii decreases
– the number of protons in the nucleus increases
» increases the strength of the positive nucleus
and pulls electrons closer to it
Rb
• IONIC RADII
– decreases across periods for same reason
as atomic radii (nucleus becomes stronger)
– alkali metals
• cations are smaller that the parent atom
– have lost an electron (actually, lost an entire level)
– therefore have fewer electrons than protons
+
Li
0.152 nm
3e and 3p
Li + ,
0.078 nm
2e and 3 p
forming a
cation
• radii still increases downwards as more levels are added on
– halogens
• anions are larger than parent atom
– have gained an electron to achieve noble gas
configuration
F 0.064 nm
9e- and 9p+
F- 0.133 nm
10 e- and 9 p+
forming
an
anion
• radii still increases downwards as more levels
are added on
– IONIZATION ENERGY
• 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
• increases across a period
– extra electrons are just filling up the same level
– the nucleus is becoming more powerful and
therefore the electrostatic force increases
making it harder to remove an electron
12.1.1
• Evidence for levels and sub-levels
– First ionization energy
• electrons are harder to remove…
– when there are more protons to attract them
– a sub-level (s,p,d,f) is completely filled
– a sub-level (s,p,d,f) are half filled
12.1.1
• Evidence for levels and sub-levels
– successive (1st, 2nd, 3rd) ionization energy
• as more electrons are removed, the electrostatic pull
of the protons holds the remaining electrons closer
• therefore, more energy is required to remove them
(even have to use a logarithmic scale to show this)
• large “jumps” are when the electrons are being
removed from the next, lower level that are much
closer to the nucleus
starting to
remove the 3s
sub-level
starting to remove
the 2s sub-level
starting
to remove
the 1s
sub-level
starting to
remove the 2p
sub-level
starting to remove the 3p sub-level
4s1 removed
– ELECTRONEGATIVITY
• 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 valence electrons (-) are shielded because of core
electrons (-) interfering with the nucleus’ (+) hold on
valence electrons
• as you go across a period
– electronegativity increases
• the atoms become smaller so the positive nucleus
can hold onto the electrons better
– MELTING POINT
• group 1 (alkali metals)
– decreases as “sea of negative
electrons” are farther away
from the positive metal ions
• group 7 (halogens)
Element
Melting
Point (K)
Li
453
Na
370
K
336
Rb
312
Cs
301
Fr
295
– increases downwards 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
decreases
• across the table (period 3)
– from left to right
• 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 can attract electrons
better
» (see next slide)
1- charge
least reactive
most reactive
• halogens
– diatomic molecules such as F2, Cl2, Br2, I2
» can react with halide ions (Cl -, Br -, and I -)
» the single bond is broken and each atom can gain one
electron to form halide ions (F1-, Cl1-, Br1-, I1-)
» the most reactive ends up as an ion (1- charge) and is not
visible (molecules F2, 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
formation of I2
due
to
due
to
Br2 no reaction
no reaction
turns
brown
formation of I2
I2 no reaction
no reaction
no reaction
– periods
• from left to right in period 3
– metals…metaloids…nonmetals
– 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