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Periodic Trends Reading
Periodic Trends
The arrangement of the periodic table reveals trends in the properties of the elements. A TREND is a
predictable change in a particular direction. For example, thpere is a trend in the reactivity of the alkali metals
as you move down Group 1. Each of the alkali metals react with water. However, the reactivity of the alkali
metals varies. At the top of Group 1, lithium is the least reactive, sodium is more reactive, and potassium is
still more reactive. In other words, there is a trend toward great reactivity as you move down the alkali metals
in Group 1.
Understanding a trend among the elements allows you to make predictions about the chemical behaviour of
the elements. These trends in properties of the elements in a group or period can be explained in terms of
electron organization. These trends can be predicted using the periodic table and can be explained and
understood by analyzing the electron configurations of the elements.
Elements tend to gain or lose valence electrons to achieve stable octet formation. Stable octets are seen in the
inert gases, or noble gases, of Group 18 of the periodic table. In addition to this activity, there are two other
important trends. First, electrons are added one at a time moving from left to right across a period. As this
happens, the electrons of the outermost shell experience increasingly strong nuclear attraction, so the
electrons become closer to the nucleus and more tightly bound to it. Second, moving down a column in the
periodic table, the outermost electrons become less tightly bound to the nucleus. This happens because the
number of filled principal energy levels increases downward within each group. These trends explain the
periodicity observed in the elemental properties of atomic radius, reactivity, ionization energy, electron affinity,
and electronegativity.
Atomic Radius
The exact size of an atom is hard to determine. An atom’s size depends on the volume occupied by the
electrons around the nucleus, and the electrons do not move in well-defined paths. Rather, the volume the
electrons occupy is though of as an electron cloud, with no clear-cut edges. In addition, the physical and
chemical state of an atom can change the size of an electron cloud. Instead, to measure an atom’s size, we
measure the atomic radius. This is the length that is half the distance between the nuclei (plural of nucleus) of
two bonded atoms. Generally, the atomic radius decreases across a period from left to right and increases
down a given group. The atoms with the largest atomic radii are located in Group I and at the bottom of groups.
Moving from left to right across a period, electrons are added one at a time to the outer energy shell. Since the
number of protons is also increasing, the nuclear charge (total charge of all the protons in the nucleus) increases
across a period. This causes the atomic radius to decrease.
Moving down a group in the periodic table, the number of electrons and filled electron shells increases, but the
number of valence electrons remains the same. The outermost electrons in a group are farther from the
nucleus as the number of filled energy shells increases. Therefore, the atomic radii increase.
Atomic Radius
1. Define atomic radius
2. How do atomic radii vary as you go from left to right and from top to bottom among the
elements in a periodic table? Why?
3. Of carbon, oxygen, and lithium, which has the smallest atomic radius? Why?
4. Of the elements gallium, potassium, and bromine, which one has the largest atomic radius?
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Periodic Trends Reading
5. Which atom in the pair, sulfur and tellurium, has the smallest atom radius? Why?
6. Which atom in the pair magnesium and barium has the larger atomic radius?
7. The elements with the largest atomic radii are found in which corner of periodic table
A. lower left-hand corner
C. lower right-hand corner
B. upper right-hand corner
D. upper left-hand corner
8. Given the representation of a chlorine atom,
which circle might represent an atom of Sulphur?
9. Using the above diagram of a chlorine atom,
which circle might represent an atom of fluorine?
10. Explain why nitrogen is a smaller atom than both boron and phosphorous
Reactivity
Reactivity refers to how likely or vigorously an atom is to react with other substances. This is usually
determined by how easily electrons can be removed (ionization energy) and how badly they want to
take other atom’s electrons (electron affinity) because it is the transfer/interaction of electrons that is
the basis of chemical reactions. Metals and non-metals each have their own trends.
When we look at the metals, we see that the reactivity increases as you move down the group.
However, as we move across a period, the reactivity decreases as you move from left to right. The
farther to the left and down the periodic table that you go, the easier it is for electrons to be given or
taken away, resulting in higher reactivity. This is a direct result of the nuclear charge and the
attraction to the outermost electrons.
The non-metals, however, show different trends. As you move from left to right in the period, we see
that the reactivity increases. On the other hand, as you move down a group in the nonmetals, the
reactivity decreases. The farther right and up you go on the periodic table, the higher the electron
affinity, resulting in a more vigorous exchange of electrons. Where metals will easily give up an
electron to gain stability, non-metals want to gain electrons.
Reactivity
1. Define reactivity
2. What are the distinguishing properties of metals and non-metals? Give three example of each.
3. In reactions, do alkaline earth metals gain or lose electrons (how many)? Explain.
4. Explain the group trend for reactivity of metals?
5. The halogen group wants to gain one electron, and the alkali metal group wants to
lose one electron. Explain why this is.
6. Why is it easier for oxygen to gain electrons than it is for selenium?
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Periodic Trends Reading
Ionization Energy
When atoms have equal amount of protons and electrons, they are electrically neutral. But when
enough energy is added, individual electrons can be removed from an atom. The neutral atom then
becomes a positively charged ion. The ionization energy, or ionization potential, is the energy required
to completely remove an electron from a gaseous atom or ion. The closer and more tightly bound an
electron is to the nucleus, the more difficult it will be to remove, and the higher its ionization energy
will be. The first ionization energy is the energy required to remove one electron from the parent atom.
The second ionization energy is the energy required to remove a second valence electron from the +1
ion to form the +2 ion, and so on. Successive ionization energies increase. The second ionization
energy is always greater than the first ionization energy.
Ionization energy tends to decrease down a group. Each element has more occupied energy levels
than the one above it does. Therefore, the outermost electrons are furthest from the nucleus near the
bottom of a group. Similarly, as you move down a group, each successive element contains more
electrons in the energy levels between the nucleus and the outermost electrons. These inner
electrons cause the outermost electrons to be held less tightly to the nucleus.
Ionization energy, however, tends to increase as you move from left to right across a period. From
one element to the next in a period, the number of protons and electrons increase by one each. The
additional proton increases the total charge of the nucleus (nuclear charge). The additional electron is
added to the same outer energy level as the previous element in the period. However, a higher
nuclear charge attracts the outer electrons more strongly. Therefore, they are held onto “tighter” and
need more energy to be removed.
Ionization Energy
7. Define ionization energy.
8. How does ionization energy vary as you go from left to right across the periodic table?
9. How does ionization energy vary as you go from top to bottom on the periodic table?
10. Of the elements calcium, beryllium or magnesium, which has the highest ionization
energy? Why?
11. Which atom in each of the following pairs has higher ionization energy?
a. S or P
b. Rb or Sr
c. Mg or Al
d. O or S
e. Mg or Ba
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Electron Affinity
Electron affinity reflects the ability of an atom to accept an electron. It is the energy change that
occurs when an electron is added to a gaseous atom. Atoms with stronger effective nuclear charge
have greater electron affinity. Some generalizations can be made about the electron affinities of
certain groups in the periodic table. The Group 2 elements, the alkaline earths, have low electron
affinity values. These elements are relatively stable because they have filled s subshells. Group 17
elements, the halogens, have high electron affinities because the addition of an electron to an atom
results in a completely filled shell. Group 18 elements, noble gases, have electron affinities near zero,
since each atom possesses a stable octet and will not accept an electron readily. Elements of other
groups have low electron affinities.
Electron Affinity
1. What is meant by electron affinity?
2. How does electron affinity vary as you go from left to right across the periodic table?
3. How does electron affinity vary as you go down a group in the periodic table?
4. Which element should have the higher electron affinity, selenium or tellurium? Why?
5. Which element should have the higher electron affinity, arsenic or antimony? Why?
Electronegativity
Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond. You
can think of it as how evenly an electron will be shared between the atoms. The more electronegative
an atom is, the more it will want to hang on to the electron (it will “hog” the electron). The higher the
electronegativity of an atom, the greater its attraction for bonding electrons. Electronegativity is related
to ionization energy. Electrons with low ionization energies have low electronegativites because their
nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have
high electronegativities due to the strong pull exerted on electrons by the nucleus. In a group, the
electronegativity decreases as atomic number increases, as a result of increased distance between
the valence electron and nucleus (greater atomic radius). An example of an electropositive (i.e., low
electronegativity) element is cesium; an example of a highly electronegative element is fluorine.
Electronegativity
6. Define electronegativity.
7. How does electronegativity vary as you go from left to right across the periodic table?
8. How does electronegativity vary as you go from top to bottom on the periodic table?
9.
Which element is the most electronegative? Why?
10. Which element is the least electronegative (that means the most electropositive)? Why?
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