Periodic table and Trends

Week No. 2-1 CHM61-101 Chemistry I Walailak University

Periodic table and Trends

Assoc. Prof. Dr. Phimphaka Harding
School of Science

Lesson objectives

• Define and write equations for ionisation energy.
• Explain how ionisation energy data provides evidence for electron structure.
• Classify elements as s, p, d or f block according to its position in the Periodic Table, which is determined by its proton number.

Learning outcomes Keywords

Students should be able to: Ionisation energies (IEs), Trends in periodic table, Group and Row, atomic
• explain how data from ionisation energies provides evidence for radius, atomic size, Electron shielding, Electronegativity, Electron affinity

electron structure.

Learning activity with opportunity to develop skills

• Students research values of first ionisation energies for elements Z=1–36 and plot them on a graph and then explain trends.
• Students write explanations for trends in ionisation energies down a group and across a period.
• Students determine which Group an element is in using successive ionisation energy data.

Periodicity

The periodic table arranges all chemical elements in special ways.
Different types of chemical bonding, and patterns and trends can be
observed in their arrangement.

All the chemical elements are arranged in the periodic table in
horizontal rows (periods) in order of increasing atomic number and
also in vertical columns (groups).
Elements in the same group have similar reactivities resulting from
a common number of outer electrons.
This allows chemists to make predictions about:
• reactivity
• physical properties
• type of bonding

Within the first 20 elements there are various different types of
bonding displayed, moving from metallic to non-metallic across the
periodic table.

Metallic bonding

Metallic bonding occurs between the atoms of metal elements -
Lithium, Berrylium, Sodium, Magnesium, Aluminium and Calcium.
The outer electrons are delocalised (free to move).
This produces an electrostatic force of attraction between the
positive metal ions and the negative delocalised electrons.
This delocalised 'sea of electrons' is responsible for metal elements
being able to conduct electricity.

Covalent molecules

Discrete covalent molecules are small groups of atoms held
together by strong covalent bonds inside the molecule and weak
intermolcular forces between the molecules.

The covalent bond itself is a shared pair of electrons
electrostatically attracted to the positive nuclei of two non-metal
atoms. The atoms achieve a stable outer electron arrangement (a
noble gas arrangement) by sharing electrons.
Most of the discrete covalent molecules are diatomic elements:
• Hydrogen (H2)
• Nitrogen (N2)
• Oxygen (O2)
• Fluorine (F2)
• Chlorine (Cl2)

There are also some larger covalent molecular elements:
• Phosphorous (P4)
• Sulfur (S8)
• Fullerenes (C60)

P4 S8 C60

Covalent network

Covalent networks are large, rigid three-dimensional arrangements
of atoms held together by strong covalent bonds. Boron, Carbon
and Silicon all form covalent networks.

They have high melting points because they only contain strong
bonds. Examples include carbon in the forms of diamond and
graphite.

CPMaotvotaenlreanntsotamnneidctwetrloernmkdesnitnsthe periodic table

Group 0 elements (the noble gases) including Helium, Neon and
Argon, exist as single, unattached particles. They are stable atoms.
They have fuller outer energy levels so they do not usually form
molecules with other atoms.

They have low melting and boiling points as they are easily
separated by overcoming the weak forces of attraction between
the atoms.

Summary of bonding

Structure Density Melting Point Conduction
Low Low Non-conductor
Monatomic Low
(noble gas) Low Non-conductor

Covalent molecular
(N2, P4, S8)

Covalent network Very high Very high Non-conductor
(except graphite)

Metallic lattice High High Conductor
(all metals)

BPalottcekrsnosfatnhde tPreernioddsiicnTtahbeleperiodic table

Periods and Blocks The fifth period is similar to the fourth. After the 6s sublevel fills, the 4f sublevel with its 14
There are seven horizontal rows of the periodic table, called periods. The electrons fills. This is followed by the 5d and the 6p. The total number of elements in the
length of each period is determined by the number of electrons that are sixth period is 32. The later elements in the seventh period are still being created. So while
capable of occupying the sublevels that fill during that period, as seen in there are a possible of 32 elements in the period, the current number is slightly less.
the Table below.
The period to which a given element belongs can easily be determined from its electron
Period Number of elements in period Sublevels in order of fill configuration. For example, consider the element nickel (Ni). Its electron configuration is
1 2 1s [Ar]3d84s2. The highest occupied principal energy level is the fourth, indicated by the 4 in
2 8 2s2p the 4s2 portion of the configuration. Therefore, nickel can be found in the fourth period of
3 8 3s3p the periodic table.
4 18 4s3d4p
5 18 5s4d5p Based on electron configurations, the periodic table can be divided into blocks denoting
6 32 6s4f5d6p which sublevel is in the process of being filled. The s, p, d, and f blocks are illustrated below.
7 32 7s5f6d7p
The figure also illustrates how the d sublevel is always one principal level behind the period
Recall that the four different sublevels each consist of a different number in which that sublevel occurs. In other words, the 3d sublevel fills during the fourth period.
of orbitals. The s sublevel has one orbital, the p sublevel has three The f sublevel is always two levels behind. The 4f sublevel belongs to the sixth period
orbitals, the d sublevel has five orbitals, and the f sublevel has seven
orbitals. In the first period, only the 1s sublevel is being filled. Since all
orbitals can hold two electrons, the entire first period consists of just two
elements. In the second period, the 2s sublevel, with two electrons, and
the 2p sublevel with six electrons, are being filled. Consequently, the
second period contains eight elements. The third period is similar to the
second, filling the 3s and 3p sublevels. Notice that the 3d sublevel does
not actually fill until after the 4s sublevel. This results in the fourth period
containing 18 elements due to the additional 10 electrons that are
contributed by the d sublevel.

Summary

• The horizontal rows of the periodic table are called periods.
• The length of a period depends on how many electrons are needed to occupy the sublevels that fill the period.
• Blocks indicate which sublevel is being filled.

Final wrap-up

Review
1. What are the horizontal rows of the periodic table called?
2. Which sublevel is being filled in period 1?
3. Which sublevel is being filled in period 7?
4. How does the electron configuration of an element give information about the period it is in?
5. What block of elements has the d sublevels being filled?

PHaytdtreorgnesnanadndtrAelnkdasliiMn tehtaelsperiodic table Element Symbol Electron configuration
Hydrogen H 1s1
Some chemistry students just enjoy learning about the science, while others are Lithium Li [He]2s1
intrigued by the violent reactions that sometimes can occur. Many chemistry classes Sodium Na [Ne]3s1
have been enlivened by the demonstration of how reactive sodium is with water. In Potassium K [Ar]4s1
some instances, the demonstration has gone off safely. Unfortunately, in other Rubidium Rb [Kr]5s1
situations students and instructors have incurred serious injury due to their failure to Cesium Cs [Xe]6s1
observe proper safety precautions. Francium Fr [Rn]7s1

One value of the periodic table is the ability to make predictions about the behavior of
individual elements. By knowing which group an element is in, we can determine the
number of reactive electrons and say something about how that element will behave.

Hydrogen and Alkali Metals

The periodic table is arranged on the basis of atomic numbers (number of protons in
the nucleus). One of the valuable consequences of this arrangement is that we can
learn a lot about the electron distribution in these atoms. The colors in the table
below indicate the different groupings of atoms based on the location and number of
electrons in the atom.

If we look at Group I (red column), we see that it is labeled alkali metals. Also note the
green H above the alkali metals. All of these elements have a similar configuration of
outer-shell electrons (see Table). In each case, there is one electron in the outer orbital
and that is an s-orbital electron. Hydrogen is not an alkali metal itself, but has some
similar properties due to its simple one proton (located in the nucleus), one electron
arrangement. The lone electron exists in a s-orbital around the nucleus. For lithium,
there are two 1s electrons in an inner orbit and one 2s electron in the outer orbit. The
same pattern holds for sodium and potassium.

HPaytdtreorgnesnanadndtrAelnkdasliiMn tehtaelsperiodic table

Even an atom with a very complex electron composition such as cesium
still has the single s electron in its outer orbital (see Figure below).

This one electron is very easily removed during chemical reactions. The
group I elements react rapidly with oxygen to produce metal oxides. They
are very soft metals, which become liquid just above room temperature.
Li reacts with water to produce hydrogen gas. Sodium also reacts the same
way, just more rapidly. Potassium reacts rapidly with water producing
hydrogen gas and heat which ignites the hydrogen gas. Rubidium and
cesium react yet more vigorously and explode on contact with water.

Summary

• Group I (alkali metals and hydrogen) elements all have one electron in their outer shell. This electron is in a s orbital.
• The Group I metals are all very reactive with water.

Final wrap-up

Review
1. What group are the alkali metals and hydrogen in?
2. What is the outer shell electron configuration in this group?
3. How reactive are the alkali metals with oxygen?
4. How reactive are these metals with water?

APalktatelirnnesEaanrdthtrMenedtaslsin the periodic table

Group 2 elements are referred to as “alkaline earth” metals (tan column below). The Uses of Alkaline Earth Compounds
name “alkaline” comes from the fact that compounds of these elements form basic (pH Since magnesium burns brightly, it is used in flares and fireworks.
greater than 7) or alkaline solutions when dissolved in water. If the Group 1 elements Magnesium alloys with aluminum provide light weight and sturdy
all have one s electron in their outer orbital, we can predict that the Group 2 elements materials for airplanes, missiles, and rockets. Several antacids use
will have two electrons in that outer shell. magnesium hydroxide to neutralize excess stomach acid.
Calcium compounds are widely found in limestone, marble, and chalk.
The beryllium atom, the first element of Group 2, has an atomic number of four. The Calcium is an important constituent of cement. Other uses include
atom has the 1s shell filled as well as the 2s shell, giving a total of four electrons calcium chloride as a deicer and limestone as a white pigment in paints
(1s22s2). Note that there are two s electrons in the outer shell, a structure that is and toothpaste.
characteristic of the Group 2 elements. Barium (atomic number 56) has the same outer
shell structure of two electrons in the s orbital, even though the internal electron Strontium is widely used in fireworks and magnets. Barium compounds
structure for barium is quite complicated. can be used in paints, filler for rubber, plastic, and resins, and as a
contrast medium for X-rays. Many beryllium compounds are toxic, but
Radium (atomic number 88) has similar properties these materials have been employed in metal alloys.
to barium and is also in the Group 2 category.
However, radium is a radioactive element and is
generally under the category of radioisotopes in
addition to being an alkaline earth metal, because it
is not a stable element.

The Group 2 elements tend to be less reactive than
their Group 1 counterparts. The need to remove
two electrons in order for the material to react
means more energy is needed for electron removal.
However, these elements are reactive enough that
they do not exist in their elemental forms in nature,
but are present as compounds.

Summary

• The alkaline earth elements are in Group 2 of the periodic table.
• These elements each have two s electrons in their outer shell.
• The alkaline earth elements are less reactive than the alkali metals.

Final wrap-up

Review
1. Why are these elements known as “alkaline earth” elements?
2. How many electrons are in the outer shell of the alkaline earth elements?
3. Are the alkaline earth elements more or less reactive than the alkali metals? Explain your

answer.
4. Is radium usually considered as part of the alkaline earth category in terms of chemistry?

Explain your answer.

NPaotbtelernGsaasneds trends in the periodic table

The reactivity of an element can give us important clues as to the electron Colors of Noble Gases
configuration of that material. If an element is extremely unreactive, this suggests that The different gases glow when an electric current is passed through
the electron configuration is such that adding or removing electrons is very unlikely. them. Many of these gases are used in displays because of their
There must be a stable electron configuration that resists further reaction. chemical inertness. They are stable and will not react with other
materials in the system. Radon also will give a reddish glow, but is not
The Group VIII (new group 18) elements are essentially chemically inert (light blue used because it is radioactive and will not retain its structure as radon
column on the right). All these elements exist as monatomic gases at room for any significant length of time.
temperature. If we look at the electron configurations, we see that helium (atomic
number 2) has a full shell of two s electrons. Since there are no electrons shielding this
shell from the nucleus, these two electrons will be very difficult to remove, making
helium unreactive.

The remaining elements in the group have full outer shells consisting of two s electrons
and six p electrons for an outer shell content of eight electrons. This particular
arrangement renders the atoms fairly unreactive. This group has been referred to as
the “inert” gases, indicating that they are chemically inert, or unreactive. Another
popular term is “noble gases,” suggesting that these gases do not like to have much to
do with the other, more common materials (or that they don’t do a lot of work).

Noble Gas Compounds
In more recent years, a number of reactions using the noble gas elements have been
discovered. Although the conventional wisdom was that the complete outer shells of
these elements would not allow them to react, some scientists believed that the outer
electrons of the larger elements (such as Rn, Xe, and Kr) were far enough away from
the nucleus that they should be able to be displaced under the right set of conditions.
The first compound (XePtF6) was made with xenon in 1962. Since then, several
compounds have been made with radon, xenon, krypton, and argon. Only helium and
neon have not formed compounds at this time.

Summary

• The noble gases are in Group VIII of the periodic table.
• Helium has a full outer shell of two s electrons.
• The other gases have full outer shells of two s and six p electrons.
• Compounds have been formed with Rn, Xe, Kr, and Ar.

Final wrap-up

Review
1. What elements comprise the noble elements?
2. What state are they in at room temperature?
3. Why is helium non-reactive?
4. Why were the other noble gases believed to be non-reactive?
5. When was the first compound formed from xenon?
6. What happens when an electric current is passed through these gases?

HPaatltoegrennssand trends in the periodic table

Some elements are much more reactive than others. The Group I (red) and Group II
(tan) elements can easily lose electrons during a reaction. Elements of other groups
are much more likely to accept electrons as they react.

The elements of Group VIIA (new Group 17 - fluorine, chlorine, bromine, iodine, and
astatine) are called the halogens (tan column). The term “halogen” means “salt-
former” because these elements will readily react with alkali metal and alkaline earth
metals to form halide salts. The halogens all have the general electron configuration
ns2np5, giving them seven valence electrons. They are one electron short of having the
full outer s and p sublevel, which makes them very reactive.

Physical Properties of Halogens
As elements, chlorine and fluorine are gases at room temperature, bromine is a dark
orange liquid, and iodine is a dark purple-gray solid. Astatine is so rare that its
properties are mostly unknown. In the picture below we see chlorine gas on the left
(green), bromine solid and vapor in the middle (orange), and solid iodine (grey) on the
right. Fluorine is not shown in the picture below because it is too corrosive and will
destroy the glass container.

None of these elements are found free in nature
because of their reactivity. Fluorine is found in
combination with cations in several minerals.
Chlorine is found in table salt, in the oceans
(which are about 2% chloride ion by weight) and
in lakes such as the Great Salt Lake in Utah. Small
amounts of bromide and iodide salts can be
found in the oceans and in brine wells in several
states.

Summary

• The halogens all have seven electrons in their outer shells.
• The electron configuration in the outer shell is ns2np5.
• As the atomic number increases, the reactivity of the halogens decreases.
• Fluorine and chlorine exist as gases at room temperature, while bromine is a liquid, and iodine is a solid.

Final wrap-up

Review
1. Pick two elements that are halogens. For each, write the name, chemical symbol, and

atomic weight.
2. What does the term “halogen” mean?
3. What is the outer shell electron configuration for the halogens?
4. What is the physical state of each halogen at room temperature?
5. Where are the halogens found?

TPraatntesritniosnanMdettraelnsds in the periodic table

Many electron configurations of elements are simple and straightforward. We can look
at the outer shell and easily understand how that set of elements will react in terms of
electron gain or loss. However, there are sets of elements that are more complex in
their behavior. One such group is called the transition elements.

Transition elements are the elements that are found in Groups 3-12 (old groups IIA-IIB)
on the periodic table (salmon-colored block in the middle of the table). The term refers
to the fact that the d sublevel, which is in the process of being filled, is in a lower
principal energy level than the s sublevel filled before it. For example, the electron
configuration of scandium, the first transition element, is [Ar]3d14s2. Remember that
the configuration is reversed from the fill order – the 4s filled before the 3d begins.
Because they are all metals, the transition elements are often called the transition
metals. As a group, they display typical metallic properties and are less reactive than
the metals in Groups 1 and 2. Some of the more familiar ones are so unreactive that
they can be found in nature in their free, or uncombined state. These include platinum,
gold, and silver. Because of this unique filling order, the transition elements are often
referred to as “d-block” elements.

Compounds of many transition elements are distinctive for being widely and vividly
colored. As visible light passes through a transition metal compound dissolved in water,
the d-orbitals absorb light of various energies. The visible light of a given energy level
which is not absorbed produces a distinctly colored solution.

Summary

• The transition elements are found in groups IIIA-IIB (new groups 3-12).
• These elements are characterized by having unfilled d sublevels.
• In general, the next higher s sublevel is already filled or has one electron missing.
• Many transition element compounds are brightly colored due to the inner-level d electron transitions.

Final wrap-up

Review
1. List five different transition elements, giving their name, chemical symbol, and atomic

number.
2. What is unique about the transition elements in terms of electron configurations?
3. Why are these elements often referred to as “d-block” elements?
4. Why do many transition element compounds have bright colors?

LPantttehrannsidaensdatnrednAdcstinitdhees periodic table

We see some hidden “layers” in chemistry. As we look at the periodic table below, we All of the others have only been artificially made in the laboratory. The
see two pink boxes – one between Ba (element 56) and Hf (element 72) and the other lanthanides and actinides together are sometimes called the inner transition
between Ra (88) and Rf (104). These elements all have unfilled f-sublevels. Because of elements.
the uniqueness of the electron configurations, these elements fit into the two boxes in
the larger periodic table. Uses of Lanthanides
Lanthanides have been widely used as alloys to impart strength and hardness
As the number of elecrons in an atom increases, we begin to see some strange to metals. The main lanthanide used for this purpose is cerium, mixed with
behaviors. Due to the way the electron energy levels work, some inner levels fill after small amounts of lanthanum, neodymium, and praseodymium. These metals
one or more outer levels do. We see this in two similar groups of elements – the are also widely used in the petroleum industry for refining of crude oil into
lanthanides and the actinides. gasoline products.

The f-Block Uses of Actinides
The actinides are valuable primarily because they are radioactive. These
The first of the f sublevels to begin filling is the 4f sublevel. It fills after the 6s sublevel, elements can be used as energy sources for applications as varied as cardiac
meaning that f sublevels are two principal energy levels behind. The general electron pacemakers and generation of electrical energy for instruments on the moon.
configuration for elements in the f block is (n - 2)f1-14ns2. The seven orbitals of the f Uranium and plutonium have been employed in nuclear weapons and in
sublevel accommodate 14 electrons, so the f block is 14 elements in length. It is pulled nuclear power plants.
out of the main body of the period table and is shown at the very bottom. Because of
that, the elements of the f block do not belong to a group, being wedged in between
Groups 3 and 4. The lanthanides are the 14 elements from cerium (atomic number 58)
to lutetium (atomic number 71). The word comes from the Greek “lanthanein”
meaning “to be hidden.” The name probably arose because these elements all hide
behind one another in the periodic table. The 4f sublevel is in the process of being
filled for the lanthanides. They are all metals and are similar in reactivity to the Group 2
alkaline earth metals.

The actinides are the 14 elements from thorium (atomic number 90) to lawrencium
(atomic number 103). The 5f sublevel is in the process of being filled. The actinides are
all radioactive elements and only the first four have been found naturally on Earth.

Summary

• Lanthanides and actinides are elements with unfilled f orbitals.
• Lanthanides are all metals with reactivity similar to group 2 elements.
• Actinides are all radioactive elements.
• Lanthanides are used in optical devices (night vision goggles), petroleum refining, and alloys.
• Actinides are found primarily in applications where their radioactivity can be used to power devices such as cardiac pacemakers.

Final wrap-up

Review
1. What electron sublevel is being filled in the lanthanides?
2. What electron sublevel is being filled in the actinides?
3. What sublevel is filled just prior to the filling of this sublevel?
4. Which actinides are found naturally on earth?
5. List some uses for lanthanides.
6. List some uses for actinides.

Patterns and trends in the periodic table A comparison of the relative covalent radii of atoms is shown in the diagram
below.
Chemists observe patterns in different properties of elements as
they are arranged in the periodic table.

Covalent radius
The covalent radius (a measure of how large individual atoms are)
shows different trends if you are moving across a period or down a
group.

Across a period from left to right, the covalent radius decreases.

As you move from left to right across the periodic table, atoms have more
electrons in their outer energy level and more protons in their nucleus.
The greater attraction between the increased number of protons (increased
nuclear charge) and electrons, pulls the electrons closer together, hence the
smaller size.

As you move down a group in the periodic table, the covalent radius
increases. Atoms increase in size.

This is because of the screening effect of the filled inner electron levels.
These electrons screen or shield the outer electrons from the nuclear charge.

Paetrtioedrnics Tarnedndtrse:nAdtos minicthReapdeiursiodic table Periodic Trend
The atomic radius of atoms generally decreases from
The size of atoms is important when trying to explain the behavior left to right across a period. There are some small
of atoms or compounds. One of the ways we can express the size exceptions, such as the oxygen radius being slightly
of atoms is with the atomic radius. This data helps us understand greater than the nitrogen radius. Within a period,
why some molecules fit together and why other molecules have protons are added to the nucleus as electrons are
parts that get too crowded under certain conditions. being added to the same principal energy level.
These electrons are gradually pulled closer to the
The size of an atom is defined by the edge of its orbital. However, nucleus because of its increased positive charge.
orbital boundaries are fuzzy and in fact are variable under different Since the force of attraction between nuclei and
conditions. In order to standardize the measurement of atomic electrons increases, the size of the atoms decreases.
radii, the distance between the nuclei of two identical atoms The effect lessens as one moves further to the right
bonded together is measured. The atomic radius is defined as one- in a period because of electron-electron repulsions
half the distance between the nuclei of identical atoms that are that would otherwise cause the atom’s size to
bonded together. increase.

Atomic radii have been measured for elements. The units for Group Trend
The atomic radius of atoms generally increases from
atomic radii are picometers, equal to 10-12 meters. As an example, top to bottom within a group. As the atomic number
increases down a group, there is again an increase in
the internuclear distance between the two hydrogen atoms in an the positive nuclear charge. However, there is also
an increase in the number of occupied principle
H2 molecule is measured to be 74 pm. Therefore, the atomic energy levels. Higher principal energy levels consist
of orbitals which are larger in size than the orbitals
radius of a hydrogen atom is 74 = 37 pm. from lower energy levels. The effect of the greater
2 number of principal energy levels outweighs the
increase in nuclear charge and so atomic radius
increases down a group.

Summary

• Atomic radius is determined as the distance between the nuclei of two identical atoms bonded together.
• The atomic radius of atoms generally decreases from left to right across a period.
• The atomic radius of atoms generally increases from top to bottom within a group.

Final wrap-up

Review
1. Define “atomic radius.”
2. What are the units of measurement for atomic radius?
3. How does the atomic radius of different elements change across a period?
4. How does atomic radius change from top to bottom within a group?
5. Explain why the atomic radius of hydrogen is so much smaller than the atomic radius of

potassium.

IPoantterns and trends in the periodic table

Atoms Are Neutral Names and Symbols

The northern lights aren’t caused by atoms, because atoms are not charged Like fluoride, other negative ions usually have names ending in –ide. Positive ions,
particles. An atom always has the same number of electrons as protons. on the other hand, are just given the element name followed by the word ion. For
Electrons have an electric charge of -1 and protons have an electric charge example, when a sodium atom loses an electron, it becomes a positive sodium ion.
of +1. Therefore, the charges of an atom’s electrons and protons “cancel The charge of an ion is indicated by a plus (+) or minus sign (-), which is written to
out.” This explains why atoms are neutral in electric charge. the right of and just above the ion’s chemical symbol. For example, the fluoride ion
Q: What would happen to an atom’s charge if it were to gain extra is represented by the symbol F-, and the sodium ion is represented by the symbol
electrons? Na+. If the charge is greater than one, a number is used to indicate it. For example,
A: If an atom were to gain extra electrons, it would have more electrons iron (Fe) may lose two electrons to form an ion with a charge of plus two. This ion
than protons. This would give it a negative charge, so it would no longer be would be represented by the symbol Fe2+. This and some other common ions are
neutral. listed with their symbols in the Table below.

Atoms to Ions Cations Chemical symbol Anions Chemical symbol
Name of ion Ca2+ Name of ion Cl-
Atoms cannot only gain extra electrons. Calcium ion H+ Chloride Br-
They can also lose electrons. In either Hydrogen ion Fe2+ Bromide F-
case, they become ions. Ions are atoms Iron (II) ion Fe3+ Fluoride O2-
that have a positive or negative charge Iron (III) ion Oxide
because they have unequal numbers of
protons and electrons. If atoms lose Q: How does the iron(III) ion differ from the iron(II) ion?
electrons, they become positive ions, or A: The iron(III) ion has a charge of +3, so it has one less electron than the iron(II) ion, which
cations. If atoms gain electrons, they has a charge of +2.
become negative ions, or anions. Q: What is the charge of an oxide ion? How does its number of electrons compare to its
Consider the example of fluorine (see number of protons?
Figure below). A fluorine atom has nine A: An oxide ion has a charge of -2. It has two more electrons than protons.
protons and nine electrons, so it is
electrically neutral. If a fluorine atom
gains an electron, it becomes a fluoride
ion with an electric charge of -1.

PIoantterns and trends in the periodic table

How Ions Form Q: Why do you think atoms lose electrons to, or gain electrons from, other atoms?
A: Atoms form ions by losing or gaining electrons because it makes them more
The process in which an atom becomes an ion is called ionization. It may stable and this state takes less energy to maintain. The most stable state for an
occur when atoms are exposed to high levels of radiation. The radiation atom is to have its outermost energy level filled with the maximum possible number
may give their outer electrons enough energy to escape from the attraction of electrons. In the case of metals such as lithium, with just one electron in the
of the positive nucleus. However, most ions form when atoms transfer outermost energy level, a more stable state can be achieved by losing that one
electrons to or from other atoms or molecules. For example, sodium atoms outer electron. In the case of nonmetals such as fluorine, which has seven electrons
may transfer electrons to chlorine atoms. This forms positive sodium ions in the outermost energy level, a more stable state can be achieved by gaining one
(Na+) and negative chloride ions (Cl-). electron and filling up the outer energy level.

Properties of Ions

Ions are highly reactive, especially as gases. They usually react with ions of opposite charge to form neutral compounds. For example, positive sodium ions and negative
chloride ions react to form the neutral compound sodium chloride, commonly known as table salt. This occurs because oppositely charged ions attract each other. Ions
with the same charge, on the other hand, repel each other. Ions are also deflected by a magnetic field, as you saw in the opening image of the northern lights.

Summary

• Atoms have equal numbers of positive protons and negative electrons, so they are neutral in electric charge.
• Atoms can gain or lose electrons and become ions, which are atoms that have a positive or negative charge because they have unequal

numbers of protons and electrons.
• The process in which an atom becomes an ion is called ionization. It may occur when atoms are exposed to high levels of radiation or

when atoms transfer electrons to or from other atoms.
• Ions are reactive, attracted or repulsed by other charged particles, and deflected by a magnetic field.

Final wrap-up

Review
1. Why are atoms neutral in electric charge?
2. Define ion.
3. Compare and contrast cations and anions, and give an

example of each.
4. Describe how ions form.
5. List properties of ions.
6. The model in the illustration below represents an atom of

lithium (Li). If the lithium atom becomes an ion, which type of
ion will it be, a cation or an anion? What will be the electric
charge of this ion? What will the ion be named? What symbol
will be used to represent it?

Ionisation energy

The first ionisation energy is the energy involved in removing one
mole of electrons from one mole of atoms in the gaseous state.

The first ionisation energy of magnesium:

Mg(g) → Mg+(g) + e- 744 kJ.mol-1

The second ionisation energy is the energy required to remove a
second mole of electrons:

Mg+(g) → Mg2+(g) + e- 1460 kJ.mol-1

The third ionisation energy shows a massive increase because it
requires an electron to be removed from magnesium’s second
energy level.

Mg2+(g) → Mg3+(g) + e- 7750 kJ.mol-1

Across a period from left to right, the ionisation energy increases.

This is due to the increase in nuclear charge having a greater pull on
the electrons and therefore more energy is required to remove
electrons.

Going down a group, the ionisation energy decreases.

This is due to the shielding or screen effect of the outer electrons
from the nucleus and so the attraction is weaker and they are more
easily removed.

A comparison of the first ionisation energies of some alkali metals is shown.

Summary

• Ionization energy refers to the amount of energy needed to remove an electron from an atom.
• Ionization energy decreases as we go down a group.
• Ionization energy increases from left to right across the periodic table.

Final wrap-up

Review
1. Define “ionization energy.”
2. Do valence electrons have larger or smaller ionization energies that the inner-shell kernel

electrons?
3. What is electron shielding?
4. Describe the trends in ionization energy from left to right across the periodic table.
5. Describe the trends in ionization energy from top to bottom of a group in the periodic

table.
6. Why is the second ionization energy for lithium so much larger than the first ionization

energy?

Electron Shielding

In general, the ionization energy of an atom will increase as we There is also a shielding effect that occurs between sublevels within the same principal
move from left to right across the periodic table. There are several energy level. Specifically, an electron in the “s” sublevel is capable of shielding electrons in
exceptions to the general increase in ionization energy across a the “p” sublevel of the same principal energy level. This is because of the spherical shape of
period. The elements of Group 13 (B, Al, etc.) have lower ionization the “s” orbital. The reverse is not true – electrons in “p” orbitals do not shield electrons in “s”
energies than the elements of Group 2 (Be, Mg, etc.). This is an orbitals.
illustration of a concept called “electron shielding.” Outer electrons
are partially shielded from the attractive force of the protons in the The spherical 3s orbital exhibits a The electron being removed from an Al atom is a 3p
nucleus by inner electrons. shielding effect on the dumbbell electron, which is shielded by the two 3s electrons as
shaped 3p orbital that is of slightly well as all the inner core electrons. The electron
The shielding effect is shown by the interior electron cloud (light blue) higher energy. This reduces the being removed from a Mg atom is a 3s electron,
shielding the outer electron of interest from the full attractive force of the ionization energy of a 3p electron which is only shielded by the inner core electrons.
nucleus. A larger shielding effect results in a decrease in ionization energy. compared to a 3s electron. Since there is a greater degree of electron shielding
in the Al atom, it is slightly easier to remove the
To explain how shielding works, consider a lithium atom. It has valence electron and its ionization energy is less than
three protons and three electrons – two in the first principal that of Mg. This is despite the fact that the nucleus of
energy level and its valence electron in the second. The valence the Al atom contains one more proton than the
electron is partially shielded from the attractive force of the nucleus of the Mg atom.
nucleus by the two inner electrons. Removing that valence
electron becomes easier because of the shielding effect. There is another anomaly between Groups 15 and 16. Atoms of Group 16 (O, S, etc.) have
lower ionization energies than atoms of Group 15 (N, P, etc.). Hund’s rule is behind the
explanation. In a nitrogen atom, there are three electrons in the 2p sublevel and each is
unpaired. In an oxygen atom, there are four electrons in the 2p sublevel, so one orbital
contains a pair of electrons. It is that second electron in the orbital that is removed in the
ionization of an oxygen atom. Since electrons repel each other, it is slightly easier to remove
the electron from the paired set in the oxygen atom than it is to remove an unpaired electron
from the nitrogen atom.

Summary

• Electron shielding refers to the blocking of valence shell electron attraction by the nucleus due to the presence of inner-shell electrons.
• Electrons in an s orbital can shield p electrons at the same energy level because of the spherical shape of the s orbital.
• Electrons in paired spin configurations are slightly easier to remove than unpaired electrons.

Final wrap-up

Review
1. Define “electron shielding."
2. Why do group 13 elements have lower ionization energies than group 2 elements?
3. What influence does a larger shielding effect have on ionization energy?
4. How do s orbit electrons affect the ionization energy of a p electron in the same shell?
5. Why do group 16 atoms have lower ionization energies than the corresponding group 15

atoms?

Electron Affinity Electron affinities (in kJ/mol) of the first five periods of the representative
elements. Electron affinities are negative numbers because energy is released.
In most cases, the formation of an anion by the addition of an electron to
a neutral atom releases energy. This can be shown for the chloride ion
formation below:

Cl + e- → Cl- + energy

The energy change that occurs when a neutral atom gains an electron is
called its electron affinity. When energy is released in a chemical reaction
or process, that energy is expressed as a negative number. The figure
below shows electron affinities in kJ/mole for the representative
elements. Electron affinities are measured on atoms in the gaseous state
and are very difficult to measure accurately.

The elements of the halogen group (Group 17) gain electrons most readily,
as can be seen from their large negative electron affinities. This means
that more energy is released in the formation of a halide ion than for the
anions of any other elements. Considering electron configuration, it is
easy to see why. The outer configuration of all halogens is ns2np5. The
addition of one more electron gives the halide ions the same electron
configuration as a noble gas, which we have seen is particularly stable.

Period and group trends for electron affinities are not nearly as regular as
for ionization energy. In general, electron affinities increase (become more
negative) from left to right across a period and decrease (become less
negative) from top to bottom down a group. However, there are many
exceptions, owing in part to inherent difficulties in accurately measuring
electron affinities.

Summary

• Electron affinity is a measure of the energy released when an extra electron is added to an atom.
• Electron affinities are measured in the gaseous state.
• In general, electron affinities become more negative as we move from left to right on the periodic table.
• In general, electron affinities become less negative from top to bottom of a group.

Final wrap-up

Review
1. Define “electron affinity."
2. Does addition of an electron to a neutral atom require energy to be absorbed or

released?
3. Describe the general trend for electron affinity values moving from left to right on the

periodic table.
4. Describe the general trend for electron affinity values moving from top to bottom in a

group on the periodic table.
5. Why is more energy released in the formation of a halide ion than with other elements?

Electronegativity

Electronegativity is a measure of an atom’s attraction for the
electrons in a bond.

Across a period from left to right the electronegativity of atoms
increases.

As you move from left to right across the periodic table, atoms have
a greater nuclear charge and a smaller covalent radius. This allows
the nucleus to attract the bonding electrons more strongly.

Going down a group, the electronegativity of atoms decreases.
As you move down a group in the periodic table, atoms increase in
size, with a greater number of energy levels.

The extra energy levels and increased covalent radius keep the
bonding electrons further away from the nucleus.

This screening effect is caused by the extra energy levels and means
that atoms further down groups have less attraction for the
bonding electrons.

Both of these trends show that fluorine is highly electronegative (it
pulls a shared pair of bonding electrons towards itself).

Summary

• Electronegativity is a measure of the ability of an atom to attract the electrons when the atom is part of a compound.
• Electronegativity values generally increase from left to right across the periodic table.
• Electronegativities generally decrease from top to bottom of a group.
• The highest electronegativity value is for fluorine.

Final wrap-up

Review
1. Define “electronegativity.”
2. How does electronegativity differ from electron affinity?
3. Why are the electronegativity values of metals generally low?
4. Describe the trend in electronegativities across the periodic table.
5. Describe the trends in electronegativities in a group of the periodic table.