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SCH 3U
Unit 1: Matter, Chemical Trends, Chemical Bonds
Notes on Trends in the Periodic Table (pg. 48 – 60)
Factor
Atomic
radius
Definition
A measurement from the centre of the
nucleus to the outermost electron (often in
nm where
1 nm = 0.000 000 001 m). [This is a tough
measurement to do. It is often estimated as
½ the distance between the nuclei of two
atoms of the same element. See Figure 3 pg.
51.]
Trend
1) Decreases Right  Left
across a period.
2) Increases
Top  Bottom in a
group.
Explanation
1) As one moves from left to right in a period, one valence electron is added to
the same ring (or energy level or shell) with each subsequent atom. Because
each electron is added to the same energy level, the amount of screening
does not change. At the same time, one more proton is successively added to
the nucleus. These changes result in an increased ENC and an increased
force of attraction between the nucleus and the electrons. Consequently,
the electrons are pulled in closer to the nucleus and the atomic radius
decreases.
3Li
4Be
5B
Decreasing radius
as electrons are added
into the same energy level
while feeling more EFC from
a nucleus with more and more
positve protons. Screening stays
constant as all valence electrons
are screened by two inner
electrons....
2) As one moves from top to bottom in a group, each new element has another
ring or energy level with electrons added to it. This means the valence
electrons are further and further from the nucleus. The attractive force to
the nucleus is screened even more each time another energy level is added.
So, the further one goes down the group, the lower the ENC felt by the
valence electrons. The electrons are held less and less tightly in the atom and
the added energy levels increase its size.
Group 1
3Li period 1
increasing atomic radius
11Na period 2
Factor
Ionic
radius
ionization
energy
Definition
A measurement from the centre of the
nucleus to the outermost electron in the ion.
-
the amount of energy required to remove
an electron from an atom or an ion in its
gaseous state
- 1st ionization energy is associated with
the removal of the 1st valence electron,
2nd ionization energy with the removal of
the second, etc. Generally speaking,
ionization energies increase as more
electrons are removed.
Example:
X(g) + 1st ionization energy  X1+(g) + 1e1Where 1e1- = 1 free electron
Trend
1) Positive ions tend to be
smaller than the atoms
they came from. As you
move L—>R across a
period, the more positive
the ion becomes and the
smaller it is.
e.g. Na1+ < Mg2+ < Al3+
2) Negative ions tend to be
larger than the atoms
they came from. As you
move
L R across a period,
the more negative the
ion becomes and the
larger it is.
e.g. N3- >O2- > F11) Increases
R  L across a period.
[So, generally, ionization
energies for metals are
lower than for nonmetals.]
2) Decreases
T  B in a group
Explanation
1) Removal of an electron from a neutral atom reduces the strength of the
electron – electron repulsions. This causes the attractive force from the
nucleus for the electrons to become strong enough to pull the electrons in
more closely.
2) Adding an electron to a neutral atom increases the electron – electron
repulsions while holding the nuclear charge the same. The electron – electron
repulsions overcome the attractive force between electrons and protons
enough to make the ionic radius larger than the radius of the atom it came
from.
1) Across a period, nuclear charge increases by increments of one while each
new valence electron is added. This means the shielding provided by nonvalence electrons remains the same, so ENC increases and the atomic radius
decreases. Thus, the further to the right we go, the more strongly the
valence electrons are attracted to the nucleus and the greater the ionization
energy.
2) Going down a group, a new energy level is added with each subsequent atom,
ensuring the valence electrons are moved further and further from the
nucleus. This increases the shielding provided by non-valence electrons,
decreases the ENC (even though the number of protons in the nucleus is
increasing) and causes the atomic radius to increase. Thus, the further down
the group we go, the less strongly the valence electrons are held and the
lower the ionization energy.
Factor
Electronegativity
-
-
Definition
a derived number on a scale of 1 – 4 that
describes the relative ability of an atom ,
when bonded, to attract electrons
the higher the number, the greater the
attraction of the atom for electrons in a
bond; the electronegativity value for an
atom can change depending on the atom it is
going to bond with
Trend
1) Increases
R  L across a
period.
2) Decreases
T  B in a
group
Electron affinity
-
the energy involved (either required or
released) when an electron is accepted by
an atom in its gaseous state
Example:
X(g) + 1e1-  X1+(g) + electron affinity energy
Where 1e1- = 1 free electron
1) Increases
R  L across a
period.
2) Decreases
T  B in a
group
Chemical
reactivity
When comparing elements in the same group,
chemical reactivity refers to how quickly or
violently one element reacts with a given other
substance compared to another element. For
example, fluorine has a rapid reaction with
sodium to form a white solid. Iodine and sodium
have to be heated to react. Fluorine and iodine
are both group 17 elements. Fluorine is the
more chemically reactive.
1) For metals,
chemical
reactivity tends to
increase as one
moves T  B in a
group,
2) For non-metals,
chemical
reactivity tends to
increase as one
moves B  T in a
group.
Explanation
1) Across a period, nuclear charge increases by increments of one while each
new valence electron is added. This means the shielding provided by nonvalence electrons remains the same, so ENC increases and the atomic radius
decreases. Thus, the further to the right we go, the more strongly the
valence electrons are attracted to the nucleus and the more likely the atom
will pull electrons toward itself in a bond, so the greater the
electronegativity value.
2) Going down a group, a new energy level is added with each subsequent atom,
ensuring the valence electrons are moved further and further from the
nucleus. This increases the shielding provided by non-valence electrons,
decreases the ENC (even though the number of protons in the nucleus is
increasing) and causes the atomic radius to increase. Thus, the further down
the group we go, the less strongly the valence electrons are held and the
less likely the atom will pull electrons toward itself in a bond, so the lower
the electronegativity value.
1) Across a period, nuclear charge increases by increments of one while each
new valence electron is added. This means the shielding provided by nonvalence electrons remains the same, so ENC increases and the atomic radius
decreases. Thus, the further to the right we go, the more strongly the
valence electrons are attracted to the nucleus so (generally) the more
(electron affinity) energy is released.
2) Going down a group, a new energy level is added with each subsequent atom,
ensuring the valence electrons are moved further and further from the
nucleus. This increases the shielding provided by non-valence electrons,
decreases the ENC (even though the number of protons in the nucleus is
increasing) and causes the atomic radius to increase. The larger the atom,
the smaller the attractive force felt between the nucleus and its valence
electrons so adding in another valence electron releases less (electron
affinity) energy.
1) When reacting chemically, metals tend to lose one or more valence electrons
to form positive ions. Going down a group, a new energy level is added with
each subsequent atom, ensuring the valence electrons are moved further and
further from the nucleus. This increases the shielding provided by nonvalence electrons, decreases the ENC (even though the number of protons in
the nucleus is increasing) and causes the atomic radius to increase. The
larger the atom, the smaller the attractive force felt between the nucleus
and its valence electrons so the easier it becomes to lose an electron – hence
the increase in reactivity as metallic atoms get larger.
2) When reacting chemically, non-metals tend to gain one or more electrons to
form negative ions. We know that the smaller the atom, the greater the
attractive force is between the valence electrons and the nucleus, the lower
the shielding, and the greater the ENC. So – the smaller the non-metallic
atom, the more reactive it is.