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1. PERIODIC TRENDS
GENERAL CHEMISTRY 1
ROMINA D. MEDIANA
Subject Teacher

2. PERIODIC TABLE

3. DEVELOPMENT OF THE PERIODIC TABLE
•The arrangement of elements in the modern periodic
table was made possible through the efforts of several
chemists, such as; Dobereiner, John Newlands, Dmitri
Mendeleev, and Henry Moseley.
•It started with Dobereiner’s“Law of Triads”. He found
a relationship among three elements where the atomic
weight of the middle element is nearly the same as
average of the atomic weights of other two elements.
•John Newlandsarranged the elements in what is known
as the “law of octaves”. He noted that the eighth
element has similar chemical properties with the first
element.

4. DEVELOPMENT OF THE PERIODIC TABLE
•Mendeleevprepared a tabulation of elements based on equivalent
weights (atomic mass) and the regular recurrence of properties of
the elements. In a few cases, the mass and the properties did not go
the same directions.
•Henry Moseleydiscovered that each element in Mendeleev’s table
was arranged in an order such that their integral positive charge
(atomic number) increased numerically from left to right and top to
bottom.
•The present periodic tableis arranged according to increasing
atomic number which also equals the number of electrons. The
electron configuration helps to predict and explain the recurrence
of chemical and physical properties.

5. •The Periodic table characterizes the known elements in increasing order
of atomic number.
•It starts on the top LEFT hand corner with Hydrogen and continue from left
to right which then repeats in the horizontal row below the last element.
•At first glance, the periodic table may seem disorganized with only a
couple elements on the top row and a block on the last row but it is very
specific in the way that they are organized.

6. THE PERIODIC TABLE IS ORGANIZED BY
PROPERTIES
Non-metals are the complete opposite so they are:
❑poor conductors of heat
❑At different phases at room temperature
❑Nitrogen, Oxygen, and Chlorine are gases
❑Silicon and Sulfur are brittle solids
❑Bromine is liquid
Metalloids have the properties from both metals and non -
metals.
Elements themselves can be one of the following metals, nonmetals
and metalloids.
Metals are:
❑Solid at room temperature (except for mercury which will be liquid)
❑Malleable (the ability to be flattened into a sheet)
❑Ductile (the ability to be made into wire wires)
❑good conductors of heat and electricity
❑shiny in appearance

7. ✓The periodic table is a chart in which elements having
similar chemical and physical properties are grouped
together.
✓The rows are called periods.
✓The vertical columns are called groups or families
according to the similarities in their properties.
✓At present, it contains 118 elements;
✓There are 18 groups or families.
✓The International Union of Pure and Applied
Chemistry (IUPAC) refers to the columns as Groups 1-18.

8. Some groups have been given collective names.
Group 1A elements are called alkali metals;
Group 2A elements are referred to as alkaline earth metals;
Group 7A elements are called halogens;
Group 8A elements are known as noble gases.
The Group A elements are classified as representative
elements or main group elements. These elements have
unfilled or filled s and p orbitals in the highest principal
quantum number.
The Group B elements are called the transition elements where
the d subshells are being filled up. This, however, is not the
universal convention.
The two separate rows at the bottom of the periodic table are
lanthanides and the actinides. Sometimes, they are referred to
as the f-block elements.

9. ANSWER THE FOLLOWING:
1. Determine whether the following elements are metals, non-metals or metalloids,
•calcium,
•phosphorus,
•silicon
•Krypton
2. Classify which elements are considered as the main group or transition metals. If
they are transition metals, state if they are lanthanides or actinides. The elements
are:
•Magnesium,
•Lanthanide
•Uranium
•Holmium
•Selenium
3. Arrange the following elements from the lowest to highest period number:
aluminum, polonium, germanium, and antimony.

10. ANSWERS:
1. Determine whether the following elements are metals, non-metals or metalloids:
•Calcium is a metal,
•phosphorus is a non-metal
•silicon is a metalloid
•krypton is a nonmetal
2. Identify which elements are considered as the main group or transition metals. If
they are transition metals, state if they are lanthanides or actinides. The elements
are:
•Magnesium, Selenium, and Lanthanide are all main group elements.
•Uranium is a transition metal which is part of the Actinide series.
•Holmium is a transition metal as well but is part of the Lanthanides.
3. aluminum, germanium, antimony, and polonium

11. THE PERIODIC LAW
states that the physical and chemical properties of the
elements recur in a systematic and predictable way when
the elements are arranged in order of increasing atomic
number.
Many of the properties recur at intervals. When the
elements are arranged correctly, the trends in element
properties become apparent and can be used to make
predictions about unknown or unfamiliar elements, simply
based on their placement on the table.

12. PERIODIC TRENDS
•Periodic trends are specific patterns that are present in the periodic table
that illustrate different aspects of a certain element, including its size and its
electronic properties.
•Major periodic trends include:
1. Electronegativity 2. ionization energy 3. electron affinity
4. atomic radius 5. melting point 6. metallic character
7. Ionic size
•Periodic trends provide chemists with an invaluable tool to
quickly predict an element's properties. These trends exist
because of the similar atomic structure of the elements within
their respective group families or periods, and because of the
periodic nature of the elements.

13. 1. THE EFFECTIVE NUCLEAR CHARGE
Z (nuclear charge) = the number of protons in the nucleus of an
atom
Z
eff(effective nuclear charge) = the magnitude of positive
charge “experienced” by an
electron in the atom
✓Z
eff increases from left to right across a period; changes very
little down a column
Shielding occurs when an electron in a many-electron atom is
partially shielded from the positive charge of the nucleus by other
electrons in the atom.
However, core electrons (inner electrons) shield the most and are
constant across a period.

14. Effective nuclear charge, Z
eff increases as you
go from left to right across a period.

15. 2. ATOMIC RADIUS
•Atomic radius which is the distance of the electron from the
nucleus within which 95% of the electron charge density is found.
•A more specific way to get atomic radius values is to get one-half
the distance between two nuclei in adjacent atoms (the
internuclear distance) in a metal solid or in a diatomic molecule.
•The covalent radius is one-half the distance between two identical
atoms joined together by a single bond.

16. Periodic Table showing Atomic Radius Trend
❑Atomic size gradually decreases from left to right across a period of
elements.
❑This means that the nucleus attracts the electrons more strongly,
pulling the atom's shell closer to the nucleus.
❑The valence electrons are held closer towards the nucleus of the atom. As a
result, the atomic radius decreases.

❑Down a group, atomic radius increases.
❑The valence electrons occupy higher levels due to the increasing quantum
number (n). As a result, the valence electrons are further away from the
nucleus as ‘n’ increases.
❑Electron shielding prevents these outer electrons from being attracted to
the nucleus; thus, they are loosely held, and the resulting atomic radius is
large.

17. Exercises
Using the periodic table, arrange the
following atoms in order of increasing atomic
radius. Explain your reasoning.
•a. C, Li, Be
•b. As, I, S
•c. P, Si, N

18. 3. IONIC RADIUS
Cations are smaller than the atoms from which they are formed.
When a metal atom loses one or more electrons to form a positive ion, the
positive nuclear charge exceeds the negative charge of the electrons in the
resulting cation.
For isoelectronic cations, the more positive the ionic charge, the smaller the
ionic charge.
Anions are larger than the atoms from which they are formed.
When a non-metal gains one or more electrons, it forms a negative ion termed
as anion. The nuclear charge remains constant, but Zeff is reduced because of
the additional electrons. The additional electrons results in increase
repulsions among the electrons in the outer shell. This results to the tendency
of the electrons to spread out more, thus increasing the size of the anion.
For isoelectronic anions, the more negative charge, the larger is the ionic
radius.
Isoelectronic species are elements or ions that have the same, or equal
number of electrons. They have same number of electrons, but they are
different in their physical and chemical properties.

19. a. Compare the size of a neutral atom of Na and a Na+ ion. Which
is larger. Explain.
The Na atom has 11 protons attracting 11 electrons. Its
electron configuration is [Ne] 3s1. This outer electron is lost
when it forms the Na+ ion. The Na+ ion has 11 protons attracting
only 10 electrons. Therefore the electrons are pulled closer to the
nucleus. The Na atom is larger than the Na+ ion: Na > Na+
b. Compare the sizes of Na+, Mg2+, and Al3+. Arrange according
to increasing size.
Na+, Mg2+, and Al3+ are isoelectronic; that is, they all
have the same number of electrons. They have 10 electrons
outside the nucleus. But for Al3+, the 10 electrons are pulled by
13 protons; for Mg2+, the 10 electrons are attracted by 12
protons; and for Na+, the 10 electrons are pulled by only 11
protons. Therefore, the sizes of the ions increase according to:
Al3+ < Mg2+ < Na+.

20. c. Compare the size of a F atom and a F–ion. Which is larger?
F has 9 protons attracting 9 electrons. A fluoride ion, F– ion
has 9 protons attracting 10 electrons. Zeff decreases for the
fluoride ion. The fluoride ion, F–, is larger than the F atom: F– > F
d. Compare the sizes of F–, O2–, and N3–. Arrange according to
increasing size.
F–, O2–, and N3– are isoelectronic. All have 10 electrons.
However, only 7 protons are attracting the 10 electrons in the
nitride ion; 8 protons are pulling in the 10 electrons in the oxide
ion; while 9 protons are attracting the 10 electrons in the
fluoride ion. Therefore, the ionic sizes increase according to F–
> O2– > N3–. .

21. 4. IONIZATION ENERGY TRENDS
❑Ionization energy is the energy required to
remove an electron from a neutral atom in its
gaseous phase.
❑Generally, elements on the right side of the
periodic table have a higher ionization energy
because their valence shell is nearly filled.
❑Elements on the left side of the periodic table
have low ionization energies because of their
willingness to lose electrons and become
cations.

22. ❑the varying energies are referred to as the
✓ first ionization energy, the
✓second ionization energy,
✓third ionization energy, etc.
❑The first ionization energy is the energy required to
remove the outermost, or highest, energy electron,
❑the second ionization energy is the energy required to
remove any subsequent high-energy electron from a
gaseous cation, etc

23. ❑Ionization energies decrease as atomic radii increase.
This observation is affected by n (the principal quantum
number) and Zeff (based on the atomic number and shows
how many protons are seen in the atom) on the ionization
energy (I). The relationship is given by the following
equation:
❑Across a period, Zeffincreases and n (principal quantum
number) remains the same, so the ionization energy
increases.
❑Down a group, n increases and Zeffincreases slightly;
the ionization energy decreases.

24. Periodic Table Showing Ionization Energy Trend

25. 5. ELECTRONEGATIVITY TRENDS
❑Electronegativity measures an atom's tendency to attract and form bonds
with electrons.
❑the most common scale for quantifying electronegativity is the Pauling
scale, named after the chemist Linus Pauling. The numbers assigned by the
Pauling scale are dimensionless due to the qualitative nature of
electronegativity.
❑This property exists due to the electronic configuration of atoms. Most
atoms follow the octet rule (having the valence, or outer, shell comprise of 8
electrons).
❑Elements on the right side of the periodic table are more energy-efficient in
gaining electrons to create a complete valence shell of 8 electrons.
❑The nature of electronegativity is effectively described thus: the more
inclined an atom is to gain electrons, the more likely that atom will pull
electrons toward itself.

26. •From left to right across a period of elements, electronegativity
increases
•From top to bottom down a group, electronegativity decreases.
•As for the transition metals, although they have electronegativity
values, there is little variance among them across the period and up
and down a group. This is because their metallic properties affect
their ability to attract electrons as easily as the other elements.
•According to these two general trends, the most electronegative
element is fluorine, with 3.98 Pauling units.
Periodic Table showing
Electronegativity Trend

27. 6. ELECTRON AFFINITY TRENDS
❑Electron affinity is the ability of an atom to accept an electron.
- electron affinity is a quantitative measurement of the energy change
that occurs when an electron is added to a neutral gas atom.


“ The more negative the electron affinity value, the higher an atom’s
affinity for electrons.”
❑Electron affinity generally decreases down a group of elements because each atom
is larger than the atom above it. This means that an added electron is further away
from the atom's nucleus compared with its position in the smaller atom.
•With a larger distance between the negatively-charged electron and the positively-
charged nucleus, the force of attraction is relatively weaker.
Therefore, electron affinity decreases.
•Moving from left to right across a period, atoms become smaller as the forces of
attraction become stronger. This causes the electron to move closer to the nucleus,
thus increasing the electron affinity from left to right across a period.

28. Periodic Table showing Electron Affinity Trend

29. 7. MELTING POINT TRENDS
•The melting points is the amount of energy required to break
a bond(s) to change the solid phase of a substance to a liquid.
•Generally, the stronger the bond between the atoms of an
element, the more energy required to break that bond.
•Because temperature is directly proportional to energy, a high
bond dissociation energy correlates to a high temperature.
•Melting points are varied and do not generally form a
distinguishable trend across the periodic table.
•Metals generally possess a high melting point.
•Most non-metals possess low melting points.

30. Figure 7. Chart of Melting Points of Various Elements

31. 9. BOILING POINT
•The temperature at which a substance undergoes a phase transition from a liquid
to a gas.
•measures of the attractive forces between atoms or molecules. Elements, such as
metals, that have strong attractive forces have higher melting points and boiling
points than nonmetals, which have very weak forces of attraction.
•The Halogens
•The physical states of the halogens at room temperature varies from gas to liquid
to solid as one moves down the group. As a result, the melting points and boiling
points increase as one moves down the group.
•(ii)The Alkali Metals
•The alkali metals show a decrease in melting points and boiling points due to the
weaker metallic bonds between atoms as their size increase down the group.

32. The Alkali
Metals


Melting Point
(K)
Boiling Point
(K)
Physical State
at Room
Temperature
Lithium 453 1615 solid
Sodium 371 1156 solid
Potassium 336 1032 solid
Rubidium 312 961 solid
Cesium 301 944 solid

33. 8. METALLIC CHARACTER TRENDS
•The metallic character of an element can be defined as how readily
an atom can lose an electron.
•From right to left across a period, metallic character increases
because the attraction between valence electron and the nucleus is
weaker, enabling an easier loss of electrons.
•Metallic character increases as you move down a group because
the atomic size is increasing. When the atomic size increases, the
outer shells are farther away.
•The principal quantum number increases and average electron
density moves farther from nucleus. The electrons of the valence
shell have less attraction to the nucleus and, as a result, can lose
electrons more readily. This causes an increase in metallic
character.

34. Metallic character relates to the ability to lose
electrons, and nonmetallic character relates to
the ability to gain electrons.
Periodic Table of Metallic Character Trend

35. PERIODIC TRENDS: SUMMARY

36. PROBLEMS
The following series of problems reviews general understanding
of the aforementioned material.
1. Based on the periodic trends for ionization energy, which
element has the highest ionization energy?
Fluorine (F) Nitrogen (N) Helium (He)
2.) Nitrogen has a larger atomic radius than oxygen.
A.) True B.) False
3.) Which has more metallic character, Lead (Pb) or Tin (Sn)?
4.) Which element has a higher melting point: chlorine (Cl) or
bromine (Br)?
5.) Which element is more electronegative, sulfur (S) or selenium
(Se)?

37. 6) Why is the electronegativity value of most noble gases zero?
7) Arrange these atoms in order of decreasing effective nuclear
charge by the valence electrons: Si, Al, Mg, S
8) Rewrite the following list in order of decreasing electron
affinity:
fluorine (F), phosphorous (P), sulfur (S), boron (B).
9) An atom with an atomic radius smaller than that of sulfur (S) is
A.) Oxygen (O) B.) Chlorine (Cl)
C.) Calcium (Ca) D.) Lithium (Li) E.) None of the above
10) A nonmetal has a smaller ionic radius compared with a metal of
the same period.
A.) True B.) False

38. SOLUTIONS
1. Answer: C.) Helium (He)
Explanation: Helium (He) has the highest ionization
energy because, like other noble gases, helium's valence
shell is full. Therefore, helium is stable and does not
readily lose or gain electrons.
2. Answer: A.) True
Explanation: Atomic radius increases from right to left on
the periodic table. Therefore, nitrogen is larger than
oxygen.
3. Answer: Lead (Pb)
Explanation: Lead and tin share the same column.
Metallic character increases down a column. Lead is
under tin, so lead has more metallic character.

39. 4. Answer: Bromine (Br)
Explanation: In non-metals, melting point increases down a
column. Because chlorine and bromine share the same
column, bromine possesses the higher melting point.
5. Answer: Sulfur (S)
Explanation: Note that sulfur and selenium share the same
column. Electronegativity increases up a column. This
indicates that sulfur is more electronegative than selenium.

40. 6. Answer: Most noble gases have full valence shells.
Explanation: Because of their full valence electron
shell, the noble gases are extremely stable and do
not readily lose or gain electrons.
7. Answer: S > Si > Al > Mg.
Explanation: The electrons above a closed shell are
shielded by the closed shell. S has 6 electrons above
a closed shell, so each one feels the pull of 6 protons
in the nucleus.
8. Answer: Fluorine (F)>Sulfur (S)>Phosphorous
(P)>Boron (B)
Explanation: Electron affinity generally increases
from left to right and from bottom to top.

41. 9. Answer: A.) Oxygen (O)
Explanation: Periodic trends indicate that atomic radius
increases up a group and from left to right across a
period. Therefore, oxygen has a smaller atomic radius
sulfur.
10. Answer: B.) False
Explanation: The reasoning behind this lies in the fact that
a metal usually loses an electron in becoming an ion
while a non-metal gains an electron. This results in a
smaller ionic radius for the metal ion and a larger ionic
radius for the non-metal ion.