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I) Periodic Law: When the elements are arranged in order of increasing atomic number, there is a periodic
repetition of their physical and chemical properties…establishing the concept of group
chemistry (a.k.a. family chemistry)
The question we should now ask is: How do we know about the atomic number?
Well, it’s time for a story.
Henry Moseley is another one of those uber-bright folks, of whom you have never heard … But as with so many
of these magnificent brains, there is quite a marvelous story to tell.
All right, you must first understand the times, and what was, and was not, known. Around 1908, it is pretty well
established that (there are):
 negative particles called electrons
 positive particles called alpha particles
 elements which chemically react similarly to their family members.
 Arrangement of the elements, on the still-developing periodic table
in 1909 is by identifying the elements with a whole number
(Hydrogen is 1; Helium is 2 and so forth). However, this whole
number is just a placement tool, and has no real meaning or
attachment to any physical characteristic of atoms or elements. Its
use is to order the elements on the developing periodic table so that
the elements line up correctly based upon similarities in chemical
behavior (the use of the number aligns the elements in the correct
Ernest Rutherford wins the Nobel Prize in Chemistry for his work
on alpha particles, and the theory of half-life regarding radioactive
So, what are the problems? Well, as of 1908 no one knows about neutrons, protons, the nucleus, nuclear charge,
electron configuration, or valence electrons.
Additionally, chemists dislike using that “whole number” to organize the elements, but they have little choice.
That whole number system works.
Chemists and physicists are confused because the elements do not line up according to verified experimental
measurements, like their relative atomic masses. Chemists and physicists cannot seem to get the relative atomic
mass measurements completely correct. Based, upon the relative atomic mass, (not the chemistries) there are
apparent problems with the periodic table positions of; argon to potassium, cobalt to nickel, and tellurium to
iodine (Later on, this same problem is seen with thorium to protactinium.) In 1908, no one knows why.
Okay, celebrate the New Year and move onto 1909.
Ernest Rutherford as the new director of the Cavendish Laboratories, and running high on his Nobel Prize in
Chemistry, decides to attack the mystery of atomic structure. He designs his classic “gold-foil experiment”.
He gets down and dirty with the gold-foil experiment (okay, he makes his students¸ Hans Geiger and Ernst
Marsden get down and dirty….). But, Rutherford crunches the data. To quote Mel Brooks, “It’s good to be the
king”. Loads of data are collected, and after the data are collected, it takes Rutherford about 2 years to crunch
the data …. So, let a couple of New Year celebrations pass by.
By 1911, Rutherford is pretty much convinced that there is a massive “nucleus” (whatever that is) at the heart of
every atom. He also knows that the nucleus is positive in charge (because it deflected the positive alpha
particles…whatever those are … he just knows their size and charge.) This positive nuclear charge of
Rutherford’s nucleus is termed “Z”.
Rutherford is clueless as to what makes up “Z”, but he hypothesizes that for atoms of the element gold, “Z” is
approximately half the mass of a gold atom. (Spoiler Alert … If you think about it, protons and neutrons are
often close to 1:1 ratio in non-radioactive isotopes, and they are almost the same in mass, so Rutherford’s guess
at “Z” being about half the mass of the nucleus is pretty darn impressive … It turns out that “Z” is really the
number of protons….but read on)
Bohr gloms onto his teacher’s (yeah, Rutherford … again!) idea of the nucleus. Bohr begins to envision
electrons around the nucleus (whatever that is…), since Rutherford showed that the positive charges and
negative charges of the atom are not co-mingled but are separated from each other.
Bohr does his bright-line emission thing for hydrogen atoms (only) … noting that the energies of the emitted
wavelengths of electromagnetic energy are equal to wavelengths of visible light, which correlate to the quanta of
energy absorbed and released by the electrons.
It is now, 1913. (Bohr wins the Nobel Prize for Physics in 1922.)
As Bohr releases his theory, Henry Moseley takes up this idea of absorbed and released energy.
Using Bohr’s work (Moseley moves pretty darn fast, evidently…), he knows that x-rays are released when
excited electrons, move from the second PEL (the L shell) down to the first PEL (the K shell).
Moseley bombards atoms of 10 or so consecutive metal elements on the periodic table, with high-speed
electrons, which displace the K-shell electrons of the bombarded atoms. (Some sources report that he bombards
12 consecutive sources … but you get the idea….)
So, as the bombarded atoms get their “K” electrons knocked out, other electrons around the nucleus drop
towards the nucleus filling the vacancy, and release characteristic x-rays.
Moseley begins to notice that the wavelengths of the released x-rays become shorter (and thus more powerful,
due to increased frequency) with increasing strength of the nuclear charge. Moseley accounts for this regular
increase in x-ray emission, by suggesting that the larger the nuclear charge, the more tightly the inner K
electrons are bound to it, thus releasing even more energy, as electrons fall into place.
Translation: The bigger “Z”, the stronger the emitted energy. The released x-ray energy depends upon the
size of the nuclear charge.
By taking the square root of the x-ray frequency, Moseley finds that the increase in energy is constant from one
element to the next. Moseley suggests that this regular increase from element to element is caused by
something in the atom by stating that there is "a fundamental quantity which increases by regular steps as we
pass from one element to the next."
Further work showed that this “something” really was the positive charge in the nucleus. Moseley referred to this
positive charge collectively as the atomic number.
CLICK Rutherford’s “Z” and Moseley’s atomic number are now seen to be the same thing!
By 1918, Rutherford (sweet glory …leave something for others to do!), suggests that there is a subatomic
particle, equal to the magnitude of the electron’s charge, but positive and the particle is far greater in mass. He
coins the term “proton”. He further suggests that the nuclear charge is an aggregation of protons.
He proves the existence of protons, using a little trick with nitrogen gas … and by 1920; he is suggesting the
existence of the neutron, to make up for the rest of the mass of an isotope. [Chadwick (another student of
Rutherford’s), wins the 1935 Nobel Prize in Physics for his 1932 discovery of the neutron…proving Rutherford
Understanding that Z and the atomic number are related, the position of the elements of the periodic table fall
into line … in that the atomic number [the nuclear charge, a.k.a. the number of protons] IS PROVEN TO BE
THE APPROPRIATE organizational schema for the elements. That silly “whole number” actually turns out to
be all-important.
By 1914, (when he was only 26 years of age!) Moseley’s work experimentally substantiates Rutherford’s goldfoil experiment, Bohr’s work on quantized electrons and quantum theory and Rutherford’s work on the proton.
Is this guy good … or what?
He is all set to walk away with the 1916 Nobel Prize in Physics by most estimates, but in 1914 WWI erupts.
Moseley as with so many of Britain’s younger generation of men, volunteers for service. In August of 1915,
during the Battle of Gallipoli, at the age of 27, he is shot dead.
His death is seen as such a blow to scientific progress, that the British Government passes a law, forbidding its
most promising scientific minds from serving in combat, ever again. It still holds to this day.
In 1962, Niels Bohr said, "You see actually the Rutherford work [the nuclear atom] was not taken seriously. We
cannot understand today, but it was not taken seriously at all. There was no mention of it any place. The great
change came from Moseley."
Kean, S. (2011). The Disappearing Spoon. Little, Brown & Company
Now, this current unit of study is all about the trends of the periodic table. And while we will speak
frequently about those electrons …please be sure to understand that the Periodic Law (the organization
based upon Moseley’s atomic number) is what leads us towards this luxury.
These are the periodic trends
we need to know
these three trends help to account for a fourth trend
(Why are some elements more easily oxidized or reduced than others?)
II) Reasons for the periodic trends really surround the ability of an atom’s nucleus * to attract its own valence
electrons and/or the electrons of another atom.
A) This ability comes down to a concept called Coulombic Forces of attraction and repulsion.
1) Essentially, August Coulomb established the idea that: * like charges repel each other
and unlike charges attract each other.
a) Thus electrons, which are negatively charged, repel each other but attract protons.
Likewise, protons repel each other. There is a mathematical expression which relates
the quantity of charge and the distance between the charges:
F ∝ Q1Q2
…. but for our work it is reasonable to summarize this by writing that:
i) the closer oppositely charged species are, the stronger the coulombic force of
ii) the closer like charged species are, there is a stronger repulsive force.
B) Attractive Coulombic Forces: Effective Nuclear Charge (Major Reason #4)
1) Effective nuclear charge an example of an attractive coulombic force and is often described
as the pull experienced by a valence electron, for a nucleus. It is not the same as the
nuclear charge (which is equal to the number of protons)
It is the resultant force experienced by a valence electron, after the effect &/or number of
shielding (inner) electrons between the valence electron and the nucleus has been taken
into account.
2) The effective nuclear charge is related to the number of valence e- of an atom, and is
increased / decreased by changes (or lack of changes), in
a) shielding effect (an example of a repulsive coulombic forces, between like-charge e-)
b) the position of the valence e- from the nucleus (due to the number of PEL and
shielding effect). This position can affect the energy of the valence e- … The
weaker the ENC, the greater the energy of the valence e-.
and possibly,
c) valence crowding when extra e- are gained (This could explain why nonmetal anions can be
oxidized back to nonmetal atoms)
What Happens: As a trend:
It becomes more & more difficult to remove an e- (to be oxidized),
conversely … gaining an e- becomes more probable
It becomes
easier and
easier to lose
an electron,
(to become
more and
Explanation of Why the Trend Happens:
The effective nuclear charge experienced by valence electrons becomes stronger & stronger. Valence e- are
bound more tightly to the nucleus and & have lower energy. This requires higher ionization energies to
remove such a valence e-. This increase in enc is due largely to the number of core electrons remaining the
same (∴ a constant shielding effect), while nuclear charge increases. Moving left to right, any added electrons
go into a valence level and valence e- do not shield each other well…(thus there is no additional shielding but
nuclear charge continues to increase, causing a resultant increase in the effective nuclear charge.) This
reasoning allows us to argue that there is an increase in attractive coulombic forces as we move from left to
right along a period.
Not only are eadded to larger
and larger PELs,
but the number of
PELs (distance
from the nucleus)
is also increasing.
This causes a
lessening of the
coulombic forces
and valance e- are
less tightly bound
to the nucleus
Evidence Used To Support What Happens:
atomic radius and metallic activity decrease,
but: 1st ionization energy, electronegativity, electron affinity, and
nonmetallic activity (the tendency to become reduced), each increase
atomic radius
and metallic
activity increase;
first ionization
energy, electron
affinity &
values decrease
III)  The changes in the periodic table trends are due to either an increase or decrease in the
coulombic forces of attraction
A) For period trends, effective nuclear charge (Zeff) is overwhelmingly important.
1) Effective nuclear charge is the net positive charge (Zeff) experience by an electron in an atom.
2) Zeff is not the full nuclear charge because of shielding effect (a.k.a. screening effect)
a) PW Atkins makes a nuanced argument regarding the interpretation of shielding effect.
He writes (p. 32): “The core electrons do not “block” in the influence of the nucleus;
they simply provide additional repulsive coulombic interactions
that partly counteract the pull of nucleus.
3) Before going on you need to know that
there are really 2 types of shielding:
the shielding of the outermost
electrons by the core electrons
the shielding of the outermost
electrons by each other!
4) In short:
*Core electrons efficiently shield electrons in the outermost principal energy
level from the nuclear charge, BUT valence electrons, do NOT efficiently shield each other.
5) So Know This: Effective nuclear charge *increases from left to right across any period
of the periodic table.
a) The number of shielding core electrons remains constant in any period … but the
nuclear charge increase, thus *the effective nuclear charge increases
b) this results in greater electronegativity values, higher ionization energies and smaller
atomic radii, as we move from left to right along a period.
6) Shielding is VERY important on period trends. It is not as important in a group trend
The primary reason for the changes in atomic properties as we move down a group
is really due to the simple fact that electrons keep being added to increasingly larger
principal energy levels (n) … thus moving farther away from the nucleus.
The increase in distance from the nucleus is the dominant effect going down a
group … as the distance increases, the attractive coulombic forces decrease.
Hence, since the radius is increasing, the attractive coulombic forces are decreasing,
and affecting the electronegativity values and ionization energies.
Summary1 for Periodic Trends
The trends occur due to a strengthening / weakening in coulombic forces
moving down a group
Group Trends: These are most easily
explained by a lessening of attractive
coulombic forces, caused by
consistent increases in the size of PEL
(affecting radius and distance from the
moving left to right along a period….
Period Trends: These are most easily
explained by a strengthening of (an
increase in) effective nuclear charge, and
thus an increase in attractive coulombic
The number of PEL is a constant.
Since Coulomb’s law states that
attraction decreases as distance
increases (inverse square rule), the
valence e- of higher atomic number
atoms of a GROUP, are expected to be
at higher energies.
The above statements best explain the
lessening of ionization energy, electron
affinity, & electronegativity, yet
increases in radius, and metallic
The number of core electrons remains the
same, while nuclear charge increases.
Any added electrons are valence e- and they
do not shield each other well … Thus
shielding is essentially constant, but the
nuclear charge increases, resulting in an
increase in the effective nuclear charge.
This primary idea of increasing ENC,
explains increases in ionization energy,
electron affinity, electronegativity &
nonmetallic activity BUT decreases in radius
Note: Avoid the argument of Shielding
Effect for group trend reasoning
B) Okay, let’s take a minute and dipstick for some understanding. Use the
information on the previous pages, dealing with “what happens”, “why it
happens”, and “the evidence we use”. Answer these questions.
1) According to the information, as you move from element to element, left to
right, across Period 3 are the valence electrons being held more tightly to the
nucleus or less tightly?
*more tightly
2) Keep comparing the elements from left to right along Period 3. Based upon your
answer to #1, should it become easier or more difficult to remove one of those
valence electrons? For instance, compare Na to Cl … is it easier or will it take
more energy to remove a valence electron from Cl (as compared to Na)?
*it will take more energy (more difficult to remove), due to the tighter hold
3) Based upon your reading and the interpretation of the information, should the first ionization
energies increase, decrease or remain the same as you move from element to element, from
left to right?
*increase, since first ionization energy measures the removal of an e-
4) Now compare the quantum configurations of
Na 1s2 2s2 2p6 3s1 & Rb 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s1
The valence electron of which atom experiences the weakest effective nuclear charge? * Rb
5) Keep comparing Na and Rb. Atoms of which element have valence electrons farthest away
from the nucleus? *Rb
6) Keep comparing Na and Rb. According to the information, atoms of which element should
lose an electron more easily? *Rb
IV) In-depth look at attractive coulombic forces e.g. Effective Nuclear Charge (MR #4 …part 1)
1) The changes in the periodic table trends are due to either an increase or decrease in the
effective nuclear charge (Zeff) (coulombic forces) for valence electrons
a) Effective Nuclear Charge: the attraction an e- "feels" for a nucleus
b) Effective Nuclear Charge decreases moving down a family, resulting in a more
loosely attracted (held) electrons ...Thus, e- are more easily lost to other atoms
c) Effective Nuclear Charge increases as you move across a period, resulting in a
tighter hold a nucleus has on its electrons ... and even more e- tend to be gained....
.... reconnect this with the Pauli Exclusion Principle and why matter has volume. Dimension (volume) allows fermions like
electrons to experience a greater level of effective nuclear charge by decreasing shielding effect ....
d) The larger the effective nuclear charge (Zeff) the smaller the radius, because of the
stronger hold on electrons
2) Effective Nuclear Charge and Family Trends:
Consider a single family. As one moves from top to bottom the number of principal energy
levels increases in the atoms of each element. The hold of the nucleus for the valence
electrons lessens with each successive principal energy level
Diagram of the PEL for Every Family of Elements
eeeeea) Because Effective Nuclear Charge (attractive coulombic forces) weaken as we
go down a family …. * valence electrons are more loosely bound to the nucleus, and
have greater energy, requiring a lesser ionization energy
3) To a lesser issue, there is something called Shielding Effect.
Shielding Effect: The result of inner e- "blocking" the positive nucleus from the valence edue to interference and mutual repulsion (Shielding Effect is an application
of repulsive coulombic forces). (MR #4 …part 2)
Recall Hund’s Rule and that
a) increasing shielding effect destabilizes the valence level by increasing electron energy
b) decreasing shielding effect creates more stable electron configuration, because
electrons are at lower energy. (MR #2, and 4 combined!!!)
4) Effective Nuclear Charge and Period Trends
Analysis of the PEL of Period 2 Elements: Groups 1 – 16
a) Notice that the “space” in which electrons are to be configured does not increase
in terms of the number of principal energy levels. This is a key point to note.
What is the consequence in terms of shielding effect, of keeping the number of inner
electrons constant, across Period 2?
* Shielding Effect changes as the number of inner electrons changes. Since the
number of inner electrons is relatively constant across a period, like Period 2, the
shielding effect is constant, or relatively unchanging.
In fact, as one moves from left to right along a period of the table, an extra proton
and an extra electron are added. Creating an increase in effective nuclear charge
V) A Short History: Working to understand the periodic trends and the work of Dimitri Mendeleev
Antoine Laurent Lavoisier divided the few elements known in the 1700’s into four classes. In the early 1800's
Dobereiner, noted that similar elements often had relative atomic masses and created triads. Dobereiner’s triads were
followed by Cannizaro who determined relative atomic masses for the 60 or so recognized elements in the 1860s. Then
Newlands arranged a table with the elements given a serial number in order of their relative atomic masses, beginning
with hydrogen. This made evident that "the eighth element, starting from a given one, is a kind of
repetition of the first", which Newlands called the Law of Octaves. This is the first hint (although incomplete) of what
is now known as the Periodic Law. However, the Periodic Law would not make complete sense until Moseley’s work
of 1913.
By 1869, both Julius Lothar Meyer and Dimitri Mendeleev (Mendeleyev) constructed periodic tables independently. The
periodicity of physical properties impressed Meyer, while Mendeleev was more interested in the chemical properties.
As Mendeleev constructed his version of the periodic table, he "left blanks" between elements, he found chemically
related, but not so closely as to go next to each other. He believed that unknown elements, would "bridge" these gaps.
Mendeleev published his periodic table in 1869 and predicted the properties of the missing elements. Chemists
began to appreciate it when the discovery of elements predicted by the table took place.
Mendeleev’s work allowed him to make predictions for undiscovered elements such as germanium, gallium, scandium,
rhenium and even technetium (which really exists only due to radioactive decay)! Additionally, Mendeleev was able to correct several
elements’ values for relative atomic mass. For example, the original atomic mass for indium was 76, due to the
assumption that indium oxide had the formula of InO. This atomic mass placed indium, which has metallic properties,
among the nonmetals. Mendeleev assumed the atomic mass was incorrect and proposed that the formula for the oxide
was really In2O3. Based upon this correct formula, indium has an atomic mass of approximately 113, placing the element
among the metals. He repeated this system of correction for beryllium and uranium. Zumdahl 3rd edition p 307 and Contemporary
Chemistry 1981 p131
Mendeleev's Genius
Studied Properties
Predicted Values for Ekasilicon
Atomic Mass
Experimentally Determined
Properties of Germanium
5.5 g/cm3
5.47 g/cm3
Specific Heat
0.31 J/(gK)
0.32 J/(gK)
Melting Point
very high
Oxide Formula (oxidation #)
Oxide Density
4.7 g/cm3
4.70 g/cm3
Chloride Formula (oxidation #)
(X = +4 oxidation state)
Specific Gravity
(Ge = +4 oxidation state)
NB: While your teacher waxes nostalgically regarding Dimitri Mendeleev’s genius ... it is fair to note that he pulled some
really odd ideas out of that magnificent brain as well. For instance, (rather inexplicitly), he put hydrogen in a group
(family) with copper, silver and mercury!!!
Additionally, Dimitri Mendeleev never accepted the concept of … the electron! This giant gaff should tell you that those
years were chaotic, and illustrate, just how important both a community of thinkers, & the freedom to think, really are.
Check out:
The values for the First Ionization Energy, Electronegativity and Atomic Radius are listed for most elements on Table S .
VI) Electronegativity: a unit-less value which implies the tendency of an atom to *attract the
electrons of a different atom
(associated with Major Reason #4)
to which it is bonded. The scale runs from 0.7 to 4.0
Trends to Know:
__Ionization Energy
__Atomic Radius
__Ionic Radius
__Metallic & Nonmetallic Activity.
A) As you move left to right across a period electronegativity values * tend to increase
As you move down a group electronegativity values * tend to decrease (just like ionization energies)
1) The Pauling Scale is a scale of 0.7 to
not likely to attract an e-
3.98 (most texts round that to 4.0, which was the original upper value)
very likely to attract the e- in a bond
2) evaluates the tendency of one atom to attract the bonding electron(s) of a different atom
in a bond.
3) In an ionic bond, one species pretty much dominates the other, rather completely, by gaining
the electron fully, while the other species loses that same electron, fully. This creates
fully realized + and – charged species
4) In a covalent bond, the species are so close in this attractive ability, that no one species can
completely dominate & thus gain the bonding electrons. Rather, there’s an accommodation or
a compromise reached between the attractive and repulsive forces, creating a sharing of the
bonding electron.
This has a tremendous consequence: This sharing creates partial + and - oxidation states
which lack the magnitude of a full positive or negative charge. This partial charge is
symbolized by a charge and the small case form of the Greek letter, delta (± )
Atoms with larger electronegativity values tend to acquire the
electrons of a bond, and the atom becomes partially negative in
oxidation state.
No difference in
the ability to gain
A :
δ: X
δA :
Bonding e- are more
attracted to the more
electronegative atom.
Electronegativity Values
Why would “not applicable” be used
to represent the electronegativity for
the noble gases?
B) Given the following bonded atoms, and the above electronegative values, assign partial charges to
each atom of the molecule.
. .
. .
. .
:O:: C::O:
1) a large electronegativity value, such as fluorine’s * 4.0
indicates that the nuclei of
atoms of fluorine have a * strong tendency / ability to attract electrons of (a) bond(s)
2) a small electronegativity value, such as cesium’s 0.7 indicates the the cesium nuclei
* have relatively no ability to attract electrons of bonds ....thus atoms of Cs, will inevitably
lose all control over the bonding electron
VII) First Ionization Energy: The number of kJ required to remove 1 mol of the most loosely held electrons
from 1 mol of atoms in the gaseous phase.
Trends to Know:
__Ionization Energy
__Atomic Radius
__Ionic Radius
__Metallic & Nonmetallic Activity.
A) As you move from left to right across a period of elements, the first ionization energy tends to
* increase / become larger
1) This means that as you go across a period, it requires more and more * energy
to * remove / lose
The process of * reduction
the most loosely held electron. Thus the process of
becomes more and more difficult and less likely to occur.
is therefore more likely as you move right.
B) As you move down a family (or group) of elements, the first ionization energy tends to
* decrease
1) This means that as you go down a group, it requires * less energy
* remove (to lose)
the most loosely held electron. Thus the process of
* oxidation (ionization)
becomes * easier (quicker, more likely)
The process of * reduction (gaining an e-)
down a group.
is therefore, less likely farther
e.g. 496 kJ + 1 mole Na0  1 mole Na1+ + 1 mole e1251 kJ + 1 mole Cl0 
1 mole Cl1+ + 1 mole e-
(a larger value suggests that this oxidation is much
less likely to occur relative to sodium's)
First Ionization Values Visualized
1) A large value such as helium’s 2,372 kJ/mol means * It requires a relatively large amount
of energy to remove the valence electron(s) of 1 mol of helium atoms.
Due to such large
energy requirements, this oxidation will probably not occur naturally, on Earth.
2) A small value such as potassium’s 419 kJ/mol means * relatively little energy is needed
to have the most loosely held electrons removed. Thus, potassium atoms are more likely
than helium to be oxidized.
3) these values help to compare atoms as to which is more likely to be * oxidized
if reacted chemically.
VIII) Covalent (Atomic) Radius: Often defined as ½ the distance between two nuclei of bonded atoms of
the same element.
A) The radius of an atom can change, depending upon the atom(s) to
which it is bonded. Thus this is a rather highly generalized trend due to
differences in measurement.
1) Trend from left to right across a period: *decrease in atomic radius
Trends to Know:
__Ionization Energy
__Atomic Radius
__Ionic Radius
__Metallic & Nonmetallic Activity.
Thus: valence electrons are * attracted to their own nucleus more strongly
______________. Thus the chance for becoming * oxidized
and the chance for becoming * reduced
2) Trend from top to bottom down a family: * increase in the size of the atomic radius
due to an increasing number of PEL’s
3) It is important to grasp the trend in atomic radius, for it can be applied to trends in ionic radius.
a) For instance …. K0 has a larger atomic radius than Na0, thus K+ while smaller than K0,
the K+ is larger than Na+ … Thus if we know the atomic radius trend, the ionic radius
trend (when comparing) will hold.
Trends in Atomic Radii as well as Changes in Ionic Radii
IX) Ionic Radius: *Generally refers to the size of an ion
A) When a metal atom is oxidized and becomes a cation, the ionic radius of the cation is
than the atomic radius of the original metal atom
1) so the cation has a *smaller radius
than the metal atom
B) When a nonmetal atom is reduced and becomes an anion, the ionic radius of the anion is
* larger
than the atomic radius of the original nonmetal atom
1) the anion has a * greater radius
than the nonmetal atom
due to increased electron /electron repulsion in the valence level. There is an increase in the repulsive forces
(due to the extra e-) as the (-) electrons interact, resulting in an increase in radius. You see, greater distance
between e- helps to reduce the effect of those repulsive forces.
X) Metallic Activity: * The tendency of an atom to be oxidized (lose electrons) in a chemical reaction
A) Study Period 2 of the Periodic Table: _______________________________________
B) Study Group 14 of the Periodic Table: ________________________________________
C) High metallic activity is linked to the nucleus’ inability to attract effectively its own valence eTrends to Know:
__Ionization Energy
__Atomic Radius
__Ionic Radius
__Metallic & Nonmetallic Activity.
Thus atoms with a high metallic activity tend to * lose
or high metallic activity is associated with atoms that are easily * oxidized
This poor attraction links up with:
Lower first ionization energy
Lower electronegativity
Larger atomic radius
Losers of electrons
XI) NONmetallic Activity: * The tendency to gain electrons (become reduced)
A) High NONmetallic activity is linked to the nucleus’ tremendous ability to hold onto its own
electrons and attract electrons from other atoms.
Thus atoms with a high NONmetallic activity tend to *gain
or high NONmetallic acitivity is associated with atoms that are easily * reduced
This strong attraction links up with: GREATER first ionization energy
GREATER electronegativity
Smaller atomic radius
The following are additional ideas regarding atomic radius to help clarify, and to consider. Essentially,
the following is the "fine print", dealing with; dioxygen, dinitrogen, the noble gases
Atomic Radius Calculations for Metals and Som Non-metals Like Oxygen and Nitrogen etc…
The atomic radius is half the distance apart of two atoms of the element in its normal state. The atomic
radius of a metal depends on its co-ordination numbers, or number of nearest neighbors. For a non-metal
the atomic radius is half the distance between a pair of bonded atoms of the element. Molecules such as
oxygen and nitrogen are multiply bonded so the atomic radii of these elements are calculated from
molecules such as hydrogen peroxide (HOOH) and hydrazine (H2NNH2) which contain only single bonds.
Gases like neon, which are atomic rather than molecular and are not known to form compounds, are more
of a problem. Their radii are calculated from the structures the elements adopt at (near) absolute zero. See
the next page, for more on the noble gases.
Atomic Radius Determination For The Noble Gases
For our current work on “trends” in atomic radius, it seems like you have to ignore the noble gas at the end of
each period. Because neon and argon don't form bonds, you can only measure their van der Waals radius - a case
where the atom is pretty well "un-squashed". All the other atoms are being measured where their atomic radius
is being lessened by strong attractions. You aren't comparing like with like if you include the noble gases.
Additionally, atomic radii are called covalent radii when referring to non-metallic elements and are
called metallic radii when referring to metals. Technically, the atomic radius is one half of the equilibrium
inter-nuclear distance between two adjacent atoms (which may either bonded covalently or present in a
closely packed crystal lattice) of an element. A covalent radius is one-half the distance between nuclei of
two of the same atoms that are bonded to each other. Covalent radii for elements whose atoms cannot bond
to each another can be estimated by combining radii of those that do with the distances between unlike
atoms in various molecules. A metallic radius is one-half of the closest inter-nuclear distance in a metallic
Effective Nuclear Charge decreases as you
move down a family, bc of more PEL and
more shielding thus causing the electrons to
be held more loosely
Effective Nuclear Charge increases as you
move across a period, bc of a buildup of
opposite charges attracting in a limited space,
thus causing the electrons to be held more
tightly to the nucleus, and thus "shrinking"
the size of the atom, despite more subatomic
DIRECTIONS: Use your reference tables and answer each question by selecting the most correct response to each.
1) Think: Periods run horizontally
The elements in Period 3 all contain the same number of:
valence electrons
occupied principal energy levels
2) Think:
What do you know about nonmetals? What information does Table S give you?
Compared to atoms of metals, atoms of nonmetals generally:
have higher electronegativity values
have lower first ionization energies
conduct electrical currents more readily
lose electrons more readily
3) Compared to an atom of potassium, an atom of calcium has a
larger radius and a less active metal
larger radius and a more active metal
smaller radius and less active metal
smaller radius and more active metal
4) Think: Hey use Table S and look up values!!!
In Period 2 of the Periodic Table, which Group contains the element with the highest first
ionization energy?
1) alkali metals
2) alkaline-earth metals
3) halogens
4) noble gases
5) As the atoms of the elements in Group1 of the Periodic Table are considered from top to bottom,
the number of valence electrons in the atoms of each successive element
1) decreases
2) increases
3) remains the same
6) Elements in the Periodic Table are arranged according to their
1) atomic number
2) atomic mass
3) relative activity
4) relative size
7) Think: What term is used to describe the “attraction for electrons”? Now look up Table S.
Which of the following elements has the strongest attraction for electrons?
1) boron
2) aluminum
3) oxygen
4) sulfur
8) Think: What is meant by “nuclear charge”?
What occurs as the atomic number of the elements in Period 2 increases?
1) The nuclear charge of each successive atom decreases, and the covalent radius decreases
2) The nuclear charge of each successive atom decreases, and the covalent radius increases
3) The nuclear charge of each successive atom increases, and the covalent radius decreases
4) The nuclear charge of each successive atom increases, and the covalent radius increases
9) An element has a first ionization energy of 1314 kJ, and an electronegativity of 3.5. It
is classified as a
1) metal
2) nonmetal
3) metalloid
4) halogen
10) The properties of carbon are expected to be most similar to those of
1) boron
2) aluminum
3) silicon
4) phosphorus
11) At which location in the Periodic Table would the most active metallic element be
1) in Group 1 at the top
2) in Group 1 at the bottom
3) in Group 17 at the top
4) in Group 17 at the bottom
12) Think: Chlorine is a nonmetal.
What occurs when an atom of chlorine forms a chloride ion?
The chlorine atom gains an electron, and its radius becomes smaller
The chlorine atom gains an electron, and its radius becomes larger
The chlorine atom loses an electron, and its radius becomes smaller
The chlorine atom loses an electron, and its radius becomes larger
13) Think: What is true about a Period of the PT?
As the elements in Period 2 of the Periodic Table are considered in succession from left to right, there
is a decrease in atomic radius with increasing atomic number. This may be explained, in part, due to
the fact that the number of protons,
and the number of
and the number of
and the number of
and the number of
shells of electrons remains the same
shells of electrons increases
shells of electrons remains the same
shells of electrons increases
14) The strength of an atom's attraction for the electrons in a chemical bond is the atom's
1) electronegativity
2) ionization energy
3) heat of reaction
4) heat of formation
15) Germanium is classified as a
1) metal
2) metalloid
3) nonmetal
4) noble gas
This equation represents the formation of a
1) fluoride ion, which is smaller in radius than a fluorine atom
2) fluoride ion, which is larger in radius than a fluorine atom
3) fluorine atom, which is the same in radius as a fluoride ion
4) fluorine atom, which is larger in radius than a fluoride ion
17) When a lithium atom forms a Li+ ion, the lithium atom
1) gains a proton
2) gains an electron
3) loses a proton
4) loses an electron
18) Electronegativity is a measure of an atom's ability to
attract the electrons in the bond between the atom and another atom
repel the electrons in the bond between the atom and another atom
attract the protons of another atom
repel the protons of another atom
19) Think: What type of elements are “brittle”?
Which element is brittle in the solid phase and is a poor conductor of heat and electricity?
1) calcium
2) sulfur
3) strontium
4) copper
20) How does the size of a barium ion compare to the size of a barium atom?
The ion is smaller because it has fewer electrons
The ion is smaller because it has more electrons
The ion is larger because it has fewer electrons
The ion is smaller because it has more electrons
21) Which of the following groups in the Periodic Table contain elements so highly reactive they
are never found in the free state?
1) 1 and 2
2) 1 and 11
3) 2 and 15
4) 11 and 15
22) Think: What class of elements have high electronegativity values?
Which element at STP is a poor conductor of electricity and has a relatively high
1) Cu
2) S
3) Mg
4) Fe
23) Think: What terms of periodic trends is associated most closely with the ability to “attract electrons”? Which element is
Which of these elements in Period 3 has the least tendency to attract electrons?
1) Mg
2) Al
3) S
4) Cl
24) In which area of the Periodic Table are the elements with the strongest nonmetallic properties
1) lower left
2) upper left
3) lower right
4) upper right
25) Think: What measurement is associated with the loss of electrons? Look up information on Table S.
Which of these metals loses electrons most readily?
1) calcium
2) magnesium
3) potassium
4) sodium
26) Bullet at least 4 reasons a student might use to explain why atoms of fluorine are reduced more readily than
atoms of bromine. (That is, provide 4 reasons why fluorine is a more active nonmetal than bromine)
27) Bullet at least 4 reasons a student might use to explain why atoms of potassium are oxidized more easily
than atoms of calcium.
28) Bullet at least 4 reasons a student can use to explain why atoms of potassium are oxidized more easily than
atoms of sodium
Answers: 1 -25
1) 4
2) 1
3) 3
4) 4
5) 3 Members of the same family based upon their common number of val. e-
6) 1
7) 3
8) 3
10) 3 It is in the same family (has the same # val. e-)
13) 1
14) 1
15) 2
17) 4
18) 1
11) 2
12) 2
9) 2
16) 2
19) 2 Did you note that 3 of them are metals and only 1 is a nonmetal? Think … this matters…
20) 1 Ba is a metal and metals lose electrons.
21) 1
22) 2 It’s asking about nonmetals. Did you note that three of the choices are metals and only one is a nonmetal?
23) 1
24) 4 Excluding the noble gases, O,F,Cl are the big nonmetal hitters.
25) 3 It has the lowest first ionization energy
26) Here are a few ideas:
atoms of fluorine have a stronger effective nuclear charge than atoms of bromine
atoms of fluorine have a greater first ionization energy
atoms of fluorine have a greater electronegativity
atoms of BROMINE have more principal energy levels, (a larger atomic radius) thus a weaker hold
on valence electrons (atoms of BROMINE have a weaker effective nuclear charge)
Variations of the above apply to 27 and 28