Unit 8 ~ The Periodic Table
8-1 History and Introduction (Section
4.8)
The Russian chemist Dimitri
Medeleev (~ 1869) arranged the
elements according to weights, but
noticed some inconsistencies (Ar to
K and Co to Ni). However, his initial
periodic table demonstrated that
the properties of elements
repeated in a pattern and he was
able to make predictions of
elements that had not been
discovered.
• A more accurate and modern version of the
table was developed by the English physicist
Henry Moseley who used x-ray
crystallography to sort the elements based
upon atomic number.
• SMART!!!!
The Periodic Law:
• the physical and chemical properties of the
elements repeat in a systematic manner with
increasing atomic number.
8-2 The Chart (Section 4.8)
• Vertical columns are called “groups” or “families” and the
elements in the group display similar chemical properties
since they all have the same # of valence electrons. Four
groups have distinct names that are commonly used:
• Group 1a: the Alkali Metals
• Group 2a: the Alkaline Earth metals
• Group 7a: the Halogens
• Group 8a: the Noble Gases
Horizontal rows are called “periods”. They represent an
“energy level” or “shell”.
Label on the chart above: transition metals, inner
transition metals, rare earths, zigzag, lanthanoids,
actinoids.
transition metals
inner transition metals
rare earths
zigzag
lanthanoids
actinoids
8-3 Trends (Section 11.11)
• Size
• Across a period: A Radii decreases.
• Reason: more protons pull on electrons
which are in orbitals with similar energy (
(greater nuclear charge).
• The F of attraction between the nucleus
and the electrons increases
Key here is that the increased # of e- are in the same E level!
• Down a group: A Radii increases.
• Reason: elements have electrons in higher energy orbitals,
which are farther from nucleus. (More shells of electrons)
• Think of the Bohr model – increased E levels = Increased
number of orbitals
Ion size
• Positive ions (cations) = loss of e- 
• Smaller +
than neutral atom due to
• Less electrostatic repulsion of the electrons
with the same number of protons (F in)
• Or even one less electron shell!
• Like OMG!!!!!
Ion size continued
• Negative ions (anions) = gain of e-
• Increased size - relative to neutral atom
• Increased F repulsion (F out) with same
number of protons (same F in)
• Also electron shielding of outer e- by inner e• Like “get away from me!!!”
A graphic for you
Isoelectronic Series
• An isoelectronic series contains atoms/ions with
the same number of electrons. For example, Cl1, Ar, and K+1 all have ______
18 electrons.
• When comparing the atomic radii in such a series,
the key is to look at the number protons for
each: The greater the number of protons pulling
on the same number of electrons, the
smaller
______________
the radius!!!!!!!!!!
Ar
K+
Cl
•
_______ > _______
> _______
Practice:
• The two cations Na+1 and Al+3 are isoelectronic
Ne
with what noble gas? ______.
10
• The ions both have ______
electrons, but Na
has 11 protons and Al has 13 protons.
• Which ion will have the smaller Atomic Radii
and why?
• Al+3 will have the smallest AR because it has
the greatest #p+ s
Ionization Energy (IE)
• Ionization energy is defined as the amount of
energy required to remove one electron from
a gaseous atom or ion.
• Remember, electrons are attracted to the
positive protons in the nucleus; it therefore
requires energy (endothermic) to remove an
electron.
• The higher the IE, the “harder” it is to remove.
•
•
•
•
Example 1:
Na + 498 kJ/mol → Na+ + e-1
Na+1 + 4560 kJ/mol → Na+2 + e-1
Na+2 + 6910 kJ/mol → Na+3 + e-1
• Note how each successive electron removed
requires more energy!
• A very big jump when removing an electron
from an lower energy level (like we see
between the 1st and 2nd e- s here.
• Example 2:
• Na (at #11) 1s2 2s2 2p6 3s1
• 1st electron removed would be 3s1
•
•
•
•
Ti (at # 22) 1s2 2s2 2p6 3s23p6 4s2 3d2
4s electrons go before the 3d electrons
Please see the orbital filling chart to see why
Strange but true!
• Across a period: IE increases
• Reason: Recall that as you go across a period the size of the atom
decreases. Smaller atoms hold electrons more tightly (e- are closer to the
nucleus).
• Down a group: IE decreases.
• Reason: Valence e- in larger atoms are farther from the nucleus.
(eare not as attracted to nucleus).
Practice (using only the periodic table)
• Which will have higher IE value (E required to
remove an e-)?:
Ar, it is higher in the same group
• Ar or Xe???
Cl, it is more to the right in same period
• P or Cl????
• K or K+ ??? K+, it has already lost one e- and new E level
• Si or Se??? Trick question
You can not tell without given values
Se is more to the right, but Si is higher
Successive Ionization Energies
• The above chart displays multiple ionization
energies. The first ionization energy (IE1) is the
energy to remove the first electron. The second
ionization energy (IE2) is the energy to remove
the second electron, and so on. What numbers
stand out as significant?
• The large jump when an E level changes!!!
• Note how the IE energies increase significantly
when an electron is removed from a full
energy level (noble gas configuration).
• Na: Write the electron configuration please
• 1s22s22p63s1
• IE1 = 498 and removes the 3s1 electron
• IE2 = 4560 and removes an e- from the Ne
noble gas configuration!!!!
Mg:
•
•
•
•
•
•
1
1s22s22p63s2
IE1 = 736
e- configuration of Na
IE2 = 1445
e- configuration of Ne
IE3 = 7730
Al:
•
•
•
•
•
•
1s22s22p63s23p1
IE1 = 577
IE2 = 1815
IE3 = 2740
IE4 = 11600
A huge jump in Energy needed to pull off that
1st e- in the 2nd Energy level!!!!
Electron Affinity (EA)
• Electron affinity is defined as the energy
change associated with adding an electron to
an atom in the gas phase to produce an anion:
X(g) + e- → X-1(g)
• Recall that IE was the E required to remove an
electron...(see previous notes) – many
students get these terms confused – You may
want to make certain you do not.
• EA values are typically negative, because
energy is usually released (exothermic).
• If the resulting anion is stable, the EA will be
negative; the more stable the anion is, the
larger the negative number.
• If the resulting anion is unstable, the value for
the electron affinity will be positive (requires E
to force the e- on!!!).
• so remember that high e- affinity = large
negative value.
• Recall that IE always required E!!!
• Electron affinity generally increases from
bottom to top within a group (that is, it goes
from smaller to larger negative numbers), and
increases from left to right within a period.
• The halogens all have large negative electron
affinities, since accepting one electron
generates a stable halide ion (with a noble gas
configuration).
• Noble gases already have a full energy level;
therefore, adding an electron will cost or
require energy, giving a +EA value.
Cl + e-1 → Cl-1 + 349 kJ
•
•
•
•
EA = -349 kJ
Is the process requiring or releasing E
Process gives off energy (▲E = -349 KJ)
Notice that the ▲E value is negative
The E as a product = releasing E
= exothermic
Recall that Cl does normally form the Cl- ion
Cl “WANTS” that e- , it is willing to pay =
high e- affinity
Mg + e-1 + 230 kJ → Mg-1 EA = +230 kJ
• Process requires energy (▲E = 230 KJ)
• Notice that the ▲E is a positive value =
requires E = endothermic
• Recall that Mg does not normally become
a negative ion – you need to force the eonto the neutral atom!!! Requires E
• So what would a energy graph of this look
like?
Graphics Please…
•
pE
Cl
pE
Mg-
Cl-
Mg
• The Chlorine anion has Less potential energy than
does the chlorine atom.
• And the magnesium atom has Less potential energy
than the magnesium anion.
Say What?????
• I like to think about the dynamite parallel –
If you give dynamite some activation
energy, it gives off a great deal of potential
energy that was stored in the bonds of the
molecules. The product (what is left after
the explosion) is not reactive, has less
potential energy, and is more stable.