Download Name: Date: Period: 1. Read the following and make annotations

Name:
Date:
1. Read the following and make annotations - Atomic Radii
Period:
The sizes of atoms vary as seen in the figure on the next page. Notice that atoms get larger as we go down a group on
the periodic table and that they get smaller as we go from left to right across a period.
We can understand the increase in size that we observe as we go down a group by remembering that as the principal
energy level increases, the average distance of the electrons from the nucleus increases. So atoms get bigger as
electrons are added to larger principal energy levels.
Explaining the decrease in atomic size across a period requires a little thought about the atoms in a given row of the
periodic table. Recall that the atoms in a particular period all have their outermost electrons in a given principal energy
level. That is, the atoms in Period 1 have their outer electrons in the 1st energy level, the atoms in Period 2 have their
outermost electrons in the 2nd energy level, and so on. Because all the orbitals in a given principal energy level are
expected to be the same size, we might expect the atoms in a given period to be the same size. However, remember
that the number of protons in the nucleus increases as we move from atom to atom in the period. The resulting increase
in positive charge on the nucleus tends to pull the electrons closer to the nucleus. So instead of remaining the same size
across a period as electrons are added in a given energy level, the atoms get smaller as the electron ‘cloud’ is drawn in
by the increasing nuclear charge.
2. Write
Atomic Radius
Atomic Radius Trend
3. Practice
1. Of the elements magnesium, chlorine, sodium,
and phosphorus, which has the largest atomic
radius? Explain your answer in terms of trends
in the periodic table.
2. Of the elements calcium, beryllium, barium, and
strontium, which has the largest atomic radius?
Explain your answer in terms of trends in the
periodic table.
3. Of the elements Li, O, C, and F, identify the one
with the largest atomic radius and the one with
the smallest atomic radius
4. Of the elements Br, At, F, I, and Cl, identify the
one with the smallest atomic radius and the one
with the largest atomic radius
Name:
Date:
1. Read the following and make annotations Ionic Radii
Period:
A positive ion is known as a cation. The formation of a cation by the loss of one or more electrons always leads to a
decrease in atomic radius because the removal of the highest-energy-level electrons results in a smaller electron cloud.
Also, the remaining electrons are drawn closer to the nucleus by its unbalanced positive charge.
A negative ion is known as an anion. The formation of an anion by the addition of one or more electrons always leads to
an increase in atomic radius. This is because the total positive charge of the nucleus remains unchanged when an
electron is added to an atom or an ion. So the electrons are not drawn to the nucleus as strongly as they were before
the addition of the extra electron. The electron cloud also spreads out because of greater repulsion between the
increase number of electrons.
As they are in atoms, the outer electrons in both cations and anions are in higher energy levels as one reads down a
group. Therefore, just as there is a gradual increase of atomic radii down a group, there is also a gradual increase of ionic
radii.
2. Write
Cation
Drawing
Anion
Drawing
3. Practice Ionic Radii
1.
2.
3.
4.
5.
A positive ion is known as a
.
A negative ion is known as an
.
Metals tend to form
.
Nonmetals tend to form
.
The formation of a cation by the
of one or more electrons always leads to a
radius because the removal of a valence electron results in a smaller electron cloud.
6. The formation of an anion by the
of one or more electrons always leads to an
radius.
Circle the correct answer…
7. Ca is a (metal/nonmetal) and will form a (cation/anion).
8. Chlorine is a (metal/nonmetal) and will form a (cation/anion).
in atomic
in atomic
Name:
Date:
1. Read the following and make annotations - Electronegativity
Period:
Valence electrons hold atoms together in chemical compounds. In many compounds, the negative charge of the valence
electrons is concentrated closer to one atom than to another. This uneven concentration of charge has a significant
effect on the chemical properties of a compound. It is therefore useful to have a measure of how strongly one atom
attracts the electrons of another atom within a compound.
Electronegativity is a measure of the ability of an atom in a chemical compound to attract electron. Electronegativities
tend to increase across each period, although there are exceptions. The alkali and alkaline-earth metals are the least
electronegative elements. In compounds, their atoms have a low attraction for electrons. Nitrogen, oxygen, and the
halogens are the most electronegative elements. Their atoms attract electrons strongly in compounds.
Electronegativities tend to decrease down a group.
2. Write
Electronegativity
Electronegativity Trend
3. Practice Electronegativity
Definition:
1. Among the elements gallium, bromine, and calcium, which has the highest electronegativity? Explain why
in terms of periodic trends.
2. Identify the most- and least- electronegative groups (families) of elements in the periodic table.
Name:
1. Read: Ionization Energies
Date:
Period:
The ionization energy of an atom is the energy required to remove an electron from an individual atom in the gas phase:
As we have noted, the most characteristic chemical property of a metal atom is losing electrons to nonmetals. Another
way of saying this is to say that metals have relatively low ionization energies—a relatively small amount of energy is
needed to remove an electron from a typical metal. This is because the metal atoms want a filled outer shell and it is
easiest for the atom to lose its few valence electrons and retain its full inner shells.
Recall that metals at the bottom of a group lose electrons more easily that those at the top. In other words, ionization
energies tend to decrease in going from the top to the bottom of a group because less energy is required to remove the
electron.
In contrast to metals, nonmetals have relatively large ionization energies. Nonmetals tend to gain, not lose, electrons.
Recall that metals appear on the left side of the periodic table and nonmetals appear on the right. Thus it is not
surprising that ionization energies tend to increase from left to right across a given period on the periodic table.
In general, the elements that appear in the lower-left region of the periodic table have the lowest ionization energies
(and therefore the most chemically active metals). On the other hand, the elements with the highest ionization energies
(the most chemically active nonmetals) occur in the upper-right region of the periodic table.
2. Write:
Ionization Energy
Metals
Nonmetals
3. Practice Ionization Energy
Definition:
In each of the following pairs of elements, circle
the element that has the higher ionization
energy…
1. Ca and Ba
2. Ca and Br
3. F and K
4. Ne and Mg
Which element has the highest ionization
energy among Se, Te, O, and S? Which
group/family does it belong to?