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Periodic Trends:
Electron Configuration:
 Notice looking at the electron configuration of the elements in family 1 (Li,
Na, K etc.), they have the same number and type of valence electrons. That
is they all have 1 electron in the s orbital. They all have s1 electrons in the
outer shell.
 Family 2 atoms also have the same number of valence electrons. This time 2
electrons in the s orbital. They all have s2 electrons in the outer shell
 Don’t worry about the transition metals – too complicated for our purposes.
 Continuing on, family 13 has 3 valence electrons, family 14 has 4 valence
electrons, family 15 has 5 valence electrons, family 16 has 6 valence
electrons, family 17 has 7 valence electrons and family 18 has zero (or 8)
valence electrons.
The number of electrons in the outermost shell, explains the shared characteristics of
elements in a family.
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The noble gases have a full outermost orbital shell. There is no need to react
in order to fill any electron holes they may have. Therefore noble gases are
inert.
The halogens have 7 electrons in the outer most shell. This is a very unstable
state. They are highly reactive in order to gain an electron to become like a
the Noble gases. This gain in 1 electron gives halogens a negative charge –
eg. FThe Alkali Earth metals have 2 valence electrons. There are 2 choices here in
order to become like Noble gases – to gain 6 electrons or loose 2 electrons.
Obviously it easier to loose 2 electrons which is what family 2 does. This give
the elements a +2 charge – eg. Ca2+
The Alkali metals are highly reactive with 1 valence electron. They loose an
electron (often to the halogens) giving them a +1 charge – eg. Na+
Isoelectronic:
two atoms/ions having the same number of valence electrons. Eg. K + is
isoelectronic to Ar. All elements aim to be isoelectronic to noble gases. This is why
elements react.
Periodic Trends:
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Reactivity:
Traveling across the Periodic Table from left to right, reactivity of elements
decreases
Traveling down each family, the reactivity increases
The reason for this is explained the in the electron configuration discussed
above
Eg1. Which element is more reactive Na or K?
Eg2. Which element is more reactive P or S?
Atomic Radii:
Left to right atomic radius decreases
o Why? As you move along each period, as the atomic mass increases,
intuitively, you expect the size to increase. However, as the number
of electrons increase, so too do the number of protons. The growing
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positive charge of the nucleus and the negative charge of the electrons
are attracted to each other. These opposing but attracted forces pull
the electrons closer to the nucleus therefore making the atom more
compact and smaller.
Down a family, atomic radius increases
o Why? We just finished saying the more electrons you get the smaller
the atom. Jumping from one row to another results in a large increase
of electrons, why the opposite trend?
o From period to period, not only do we gain electrons, but we gain
entire orbitals of electrons. The added electron orbitals create an
electron-electron force of repulsion, pushing the orbitlals away from
the nucleus therefore causing the atom radius to increase. This effect
is known as the shielding effect.
Eg 3. Which atom has the greater atomic radius, Ca or Br?
Eg 4. Which atomic radius is greater K or Cs?
Eg 5. Which atom is bigger Ni or Zr?
Note shielding effect is stronger than the effect of
adding one electron at a time across a period
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Ionization Energy:
From your lab you should recall, across a period, ionization energy increases
o Why? Similar to why the atomic radius decreases. The increase in
electrons also means an increase in protons which means a greater
force of attraction as you move across a period. As the force of
attraction increases, more energy is needed to remove the electron
Down a family, ionization energy decreases
o Why? Again similar to atomic radius trend. The shielding effect
weakens the force between the valence electrons and the nucleus.
Therefore as you move down from period to period, the electrons are
being pushed away and you need less ionization energy to remove an
electron.
This is the energy needed to remove an electron from a neutral atom
Elements with low ionization energies from positive ions because it doesn’t
take much to remove the ion and produce a positive ion
Elements with high ionization energies form negative ions because it is easier
to accept an electron than to take energy to remove one.
Eg 6. Which has the higher ionization energy Ta or Po?
Eg 7. Which has the higher ionization energy Co or Ir?
Eg 8. Which has the higher ionization energy Ti or Hg?
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Electron Affinity:
Across a period electron affinity increases
Down a family, electron affinity decreases
Electron affinity is how much an atom likes it’s own electrons when it is an
ion. The more it likes it’s electron the more energy is required to remove it
(ionization enery).
Electronegativity:
Across a period electronegativity increases
Down a family, electronegativity decreases
Electronegativity is how much an atom likes it’s electrons while forming a
compound
Metallic Properties:
Across a period elements become less metallic
Down a family elements become more metallic
Periodic Table Trends Summary:
Ionization energy decreases
Electronegativity decreases
Electron Affinity decreases
Atomic Radius decreases
Metallic Properties decrease
With this in mind, try some of the following question.
a)
State the chemical family or group of the following elements:
1) Radon Noble Gases
2) Calcium Alkaline Earth Metals
3) Iron Transition Metals
4) Cesium Alkali Metals
5) Iodine Halogens
6) Zinc Transition Metals
b)
Give the symbols for two other elements in the same family as the
following
1) Na - Li, K, Rb, Cs, Fr
2) Ar - Ne, He, Kr, Xe, Rn
3) Mg - Be, Mg, Ca, Sr, Ba, Ra
4) Br - F, Cl, Br, I, At
c)
Give tow other elements in the same period as the following
1) C - Li, Be, B, N, O F, Ne
2) S - Na, Mg, Al, Si, P, S, Cl, Ar
d)
Circle the more metallic element in the given pair.
1) B or Ga
4) Br or Sn
2) Ge or S
5) Mg or P
3) As or Bi
e)
Arrange the following in order from most metallic to least metallic: P, Ca,
F, Si, Ge. Ca, Ge, Si, P, F
f)
Circle the better electrical conductor in the given pair. ***Note,
Electrical conductivity is a property of a metallic element, so the
more metallic, the more it conducts!
1) Sb or P
4) B or Al
2) K or I
5) Tl or S
3) Ge or As
6) Sb or Te
g)
What happens to the distance between the nucleus and outer most
electrons going down a chemical family?
It increases due to the shielding effect – electrons in adjacent
orbitals repel each other.
h)
What happens to the ionization energy going down a family?
It decreases
i)
What happens to the distance between the nucleus and the outermost
electrons going left to right across a period?
It decreases due to greater positive nuclear charge attracting
the electrons and therefore “pulling in” and causing atom to
get smaller
j)
What happens to the ionization energy going across the period?
Ionization energy increases
k)
Circle the element which should have a greater ionization energy.
1) Br or Cl
4) Rb or I
2) Al or Cl
5) F or Ne
3) Ne or Xe
6) Mg or Ba