Transcript Periodic Trends

Background on the Periodic Table
Dmitri Mendeleev: given credit
for Periodic Table (~1870)
-- organized Table by
increasing atomic mass
-- left spaces and predicted
properties of undiscovered
elements
Mendeleev
Henry Moseley: put elements in
order of increasing
atomic number
____________.
Moseley
Describing the Periodic Table
periodic law: the properties of elements repeat
every so often
period: horizontal row; there are 7
group (family): vertical column; there are 18
18
1
1
2
3
4
5
6
7
1314151617
2
3 4 5 6 7 8 9 101112
Regions of the Table (cont.)
metalloids (semimetals): “stair” between metals
and nonmetals (B, Si, Ge, As, Sb, Te, Po)
metals
computer chips
properties: in-between those of metals
and nonmetals; “semiconductors”
Si and Ge
computer chips
Same number of valence e– = similar properties
Li 1s2 2s1
Na 1s2 2s2 2p6 3s1
In any group, the element BELOW has one more
occupied energy level than does the element ABOVE.
Li
Na
The period that an element is in is the same as the
energy level that its valence electrons are in.
in 2nd period
in 3rd period
Li
Na
v.e– in 2nd E.L.
v.e– in 3rd E.L.
alkali metals: group 1 (except H); 1+ charge;
very reactive
alkaline earth metals: group 2; 2+ charge;
less reactive than alkalis
halogens: group 17; 1– charge; very reactive
noble gases: group 18; no charge; unreactive
lanthanides: elements 58–71
actinides: elements 90–103
contain f
orbitals
coinage metals: group 11
transition elements: groups 3–12; variable charges
main block (representative) elements: groups 1, 2,
13–18
actinides
alkali metals
alkaline earth metals
coinage metals
halogens
transition elements
metalloids
noble gases
lanthanides
main block elements
more nonmetals
hydrogen
more metals
Periodic Trends
We will be studying four trends across the Periodic Table.
Atomic radii: the distance from the nucleus to the farthest e-.
Ionic radii: the distance from the nucleus to the farthest ein that element’s ion.
Electronegativity: the propensity for an element to attract efrom another atom. It’s ability to bond.
Ionization energy:: the amount of energy required to pull one eaway from an element; to remove one electron.
Trends
Electronegativity
Ionization energy
Atomic radius
Periodic Trend
Electron shielding: the inner e- block some of the pull from the
nucleus so the valence e- don’t feel as much.
Feels the force of 5 p+
- -
+++
++ +
+
-
Feels the force of 7 p+
Effective nuclear charge is how many p+ are pulling on that e-
shielding effect: kernel e– “shield” valence e–
from attractive force of the nucleus
As we go
, shielding effect increases...
v.e–
v.e–
Li
K
tougher to
remove
-- caused by kernel and valence e–
repelling each other
easier
to
remov
e
Shielding Effect
Valence
-
+
nucleus
Kernel electrons block
the attractive force of
the nucleus from the
valence electrons
-
-
-
Electron
Shield
“kernel”
electrons
Electrons
Periodicity
there are trends in properties of elements
-- left-right AND up-down trends
atomic radius: the size of a neutral atom
…increases as we go
WHY? add a new energy
level each time
…decreases as we go
WHY? it has to do with…
coulombic attraction: attraction between (+) and (–)
Coulombic attraction depends on…
amount of charge
2+
2–
1+
1–
distance between charges
2+
2+
–
+
H
He
++
+ +
– –
– –
2–
2–
As we go
,
more Coulombic
attraction, No new
energy level, more
pull = smaller
size
Decreasing Atomic Size
Across a Period
As the attraction between the (+) nucleus and the (–) valence electrons , the
atomic size .
From left to right, size decreases because there is an increase in nuclear
charge and Effective Nuclear Charge (# protons – # core electrons).
Each valence electron is pulled by the full Effective Nuclear Charge.
Li
1s22s1
Li
Be
1s22s2
Be
B
1s22s22p1
B
Atomic Radii
IA
IIA
VA
VIA
VIIA
Li
1.52
Be
1.11
B
0.88
C
N
O
0.77
0.70
0.66
F
0.64
Na
Mg
Al
Si
P
S
Cl
1.86
1.60
1.43
1.17
1.10
1.04
0.99
K
Ca
Ga
Ge
As
Se
Br
2.31
1.97
1.22
1.22
1.21
1.17
1.14
Rb
Sr
In
Sn
Sb
Te
I
2.44
2.15
1.62
1.40
1.41
1.37
1.33
Cs
Ba
Tl
Pb
2.62
2.17
IIIA
1.71
IVA
1.75
Bi
1.46
=1 Angstrom
Relative Size of Atoms
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 350
ionic radius: the size of an ion
cations
Ca atom
anions
Ca2+ ion
Cl atom
Cl1– ion
20 p+
20 p+
17 p+
17 p+
20 e–
18 e–
17 e–
18 e–
Cl
Cl1–
Ca
Ca2+
cations are smaller
anions are larger
Formation of Cation - Cation is Smaller than Parent.
sodium atom
Na
sodium ion
Na+
ee-
e-
e-
e-
e-
ee-
e-
11p+
e-
e-
e-
e-
e-
loss of
one valence
electron
e-
11p+
ee-
eee-
e-
Formation of Anion - Anion is Bigger than Parent Atom
chloride ion
Cl1-
chlorine atom
Cl
e-
ee-
e-
e-
e-
e-
ee-
e-
gain of
one valence
electron
e-
e-
e-
e-
e-
e-
e-
e-
e-
e-
17p+
e-
ee-
e-
e-
e-
17p+
e-
e-
e-
e-
e-
e-
ee-
e-
e-
Sizes of ions: electron repulsion
Valence electrons repel each other.
• When an atom becomes an
anion (adds an electron to its
valence shell) the repulsion
between valence electrons
increases without changing
effective nuclear charge.
• Thus, F– is larger than F
- -- 9+ - - Fluorine atom
F
1s22s22p5
-
+1e-
-
- 9+ - Fluoride ion
Fluorine
F11s22s22p6
9
+
Trends in Atomic and Ionic Size
Nonmetals
Metals
Group 1
Group 13
Group 17
e
e
Li+
Li
F
64
152
60
ee
e
e
Na+
Na
95
Al
143
e
Al3+
Cl
K+
K
Cl181
e
Br
114
133
136
99
50
186
e
F-
Br195
Anions are larger than parent atoms
Cations are smaller than parent atoms
227
IA
Atomic
Radii
Li
1.52
Cations:
smaller
than parent
atoms
IIIA
IVA
VA
VIA
VIIA
Be
B
C
N
O
F
1.11
0.88
0.77
0.70
0.66
0.64
Mg
Al
Si
P
S
Cl
1.10
1.04
0.99
Na
1.86
1.60
K
2.31
Rb
2.44
Ca
1.97
Sr
2.15
1.22 1.22
In
Sn
1.62 1.40
1.21
1.17
Sb
Te
1.41
1.37
1.14
I
1.33
Cs
Ba
2.17
Tl
Pb
1.71 1.75
Bi
1.46
O21.40
S21.84
F11.36
Cl11.81
Se2-
Br1-
1.98
Te22.21
1.85
I12.16
2.62
Ionic
Radii
IIA
1.43 1.17
Ga
Li1+
Be2+
0.60
Na1+
0.31
Mg2+
Al3+
0.95
K1+
0.65
Ca2+
0.50
Ga3+
1.33
Rb1+
0.99
Sr2+
0.62
In3+
1.48
Cs1+
1.13
Ba2+
0.81
Tl3+
1.69
1.35
0.95
Ge
As
N31.71
Se
Br
Anions:
LARGER
than parent
atoms
= 1 Angstrom
Ionization Energies
• Energy is required to remove an electron from
an atom to form a cation.
• Ionization energy () is the amount of energy
needed to remove an electron from the gaseous
atom E in its ground state:
• Larger values of  mean that the electron is
more tightly bound to the atom and is harder to
remove.
• Units for ionization energies are kilojoules/mole
(kJ/mol) or electron volts (eV) - 1 eV = 96.49
kJ/mol.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Electronegativity
The ability of an
atom in a compound
to attract shared
electrons to itself.
Linus Pauling
1901 - 1994
Electronegativity:
It is a measure of the ability of an atom to attract additional
electron to it.
Group: Electronegativity decreases down a group because
the outer energy level is further away from the nucleus which
results in a weaker nuclear charge available to attract additional
electrons or less ability for an atom to attract electrons to it.
Electronegativity (Cont.):
Period: Electronegativity increases moving across a period
from left to right because the atoms in the same period have the
same number of energy levels but the number of protons
increase as you go from atom to atom across a period
increasing the attraction between the nucleus and the outer
energy level resulting in a greater ability for atoms to attract
electrons to it.
electronegativity: the tendency for
a bonded atom to
attract e– to itself
Linus Pauling quantified
the electronegativity scale.
As we go , electronegativity… decreases.
As we go
, electronegativity… increases.