1. Periodic Trends
Elemental Properties and Patterns

2. The Periodic Law
 Though he didn’t realize it at the time, Mendeleev
made an important discovery about the elements.
 Mendeleev’s repeating patterns on his Periodic Table
led to the formulation of the ‘Periodic Law’ which
states:
 When arranged by increasing atomic number, the
chemical elements display a regular and repeating
pattern of chemical and physical properties.

3. What Does the Periodic Table Tell Us?
• Recall that the blocks of the periodic table are
arranged in to LEFT TO RIGHT ROWS called a
PERIOD/SERIES and UP AND DOWN COLUMNS
called a FAMILY/GROUP.
• Each square on the periodic table is known as
an ELEMENT KEY and provides the basic
information about each element:
• CHEMICAL SYMBOL – a one or two letter
abbreviation for the name of an element.
– Some symbols are based on the Greek or Latin
name of the element.
• Example Lead = Pb from Plumbus (where we get plumber
from).
• Gold = Au from Arium (to shine like the sun).

4. • ELEMENT NAME – the name of the element.
Some names are more than two thousand years old!
• ATOMIC NUMBER – Number of protons (nucleus) or
electrons (electron cloud).
• ATOMIC MASS NUMBER – combined number of protons
and neutrons in the nucleus.

5. VALENCE ELECTRONS
 One of the most important thing about an atom is the
number of VALENCE ELECTRONS it has.
 Valence Electrons are those electrons found on the
outside edge of the atom, farthest away from the
nucleus.

6. Valence Electrons are those on
the outer edge of the atom.

7. • An atom will have from 1 to 8 valence
electrons, abbreviated as Ve- .
• Valence electrons are used to bond to
other atoms to form COMPOUNDS.

8. Valence Electrons
• Determining the number of Valence
Electrons an element has is easy:
• Look at the GROUP NUMBER the element
is in.
• Remember the GROUP NUMBER is the numbers 1 to 18 across the
tops of the vertical columns.
• IGNORE the TRANSITION ELEMENTS (groups 3-12) for
now.
• ONLY LOOK AT GROUPS 1, 2, 13, 14, 15, 16, 17 and
18.

9.  The number of Valence Electrons is from the
group number (Groups 13-18 just use the
SECOND number):
 Group 1 = 1 Ve-
 Group 2 = 2 Ve-
 Group 13 = 3 Ve-
 Group 14 = 4 Ve-
 Group 15 = 5 Ve-
 Group 16 = 6 Ve-
 Group 17 = 7 Ve-
 Group 18 = 8 Ve-
 We will learn the Transition Elements later!

10. If you can’t read this table, follow this link to display it online:

11. Grouping the Elements
• Elements are often classified by their
type:
• Metals
• Nonmetals
• Metalloids or Semi-metals

12. Metals, Nonmetals, Metalloids

13. Metals, Nonmetals, Metalloids
 There is a zig-zag or
staircase line that
divides the table.
 Metals are on the
left of the line, in
blue.
 Nonmetals are on
the right of the line,
in orange.

14. Metals, Nonmetals, Metalloids
 Elements that border
the stair case, shown
in purple are the
metalloids or semi-
metals.
 There is one
important exception.
 Aluminum is more
metallic than not.

15. Metals  Metals are:
 Lustrous
 Have a shiny surface.
 Malleable
 Can be hammered thin.
 Ductile
 Can be stretched (wire)
 Good Conductors
 Allow heat and electricity to flow
through easily.
 They are mostly solids at room temp.
 What is one exception?
 Mercury (Hg) is a LIQUID metal!

16. Nonmetals  Nonmetal properties are the basically
opposite of metals:
 Dull
 Not shiny/can’t be polished.
 Brittle
 Break easily…cannot be hammered
thin.
 Poor/non-conductors
 Limited or no flow of heat or
electricity through them.
 (also called insulators).
 Some are solid, several are gases,
and BROMINE is a liquid.

17. Metalloids  Metalloids, aka semi-metals are
just that.
 They have characteristics of both
metals and nonmetals.
 They are shiny but brittle.
 They are semiconductors:
elements that have conduction
capacities between non metals and
metals.
 What are our most important
semiconductors?
 Silicon, Germanium, and Arsenic.

18. Periodic Trends
 ATOMIC RADIUS
 The first and most important is atomic radius.
 Radius is the distance from the center of the
nucleus to the “edge” of the electron cloud.

19. Atomic Radius
 Since a cloud’s edge is difficult to define,
scientists use a defined covalent radius, or
half the distance between the nuclei of 2
bonded atoms.
 Atomic radii are usually measured in
picometers (pm = 1 x 10-9
m) or angstroms
(Å). An angstrom is 1 x 10-10
m.

20. Example of Covalent Radius
 Two Br atoms bonded together are 2.86 angstroms
apart.
 Arad = 2.86 Å = 1.43 Å
 2
2.86 Å

21. Atomic Radius
 The trend for atomic radius in a FAMILY/GROUP
(vertical column) is to go from smaller at the top to
larger at the bottom of the family.
 Why?
 With each step down the family, we add an entirely
new PEL to the electron cloud, making the atoms
larger with each step.

22. Atomic Radius
 What happens to atomic structure as we
step from left to right?
 Each step adds a proton and an electron
(and 1 or 2 neutrons).

23. Atomic Radius
 The effect is that the more positive
nucleus has a greater pull on the electron
cloud.
 The nucleus is more positive and the
electron cloud is more negative.
 The increased attraction pulls the cloud in
closer to the nucleus, making atoms
smaller as we move from left to right
across a period.

24. Overall Trend of Atomic
RadiusINCREASESTOWARDBOTTOMLEFT
DECREASES TOWARD TOP RIGHT

25. e- Shielding
 As more PELs are added to atoms, the inner levels of
electrons shield the outer electrons from the
nucleus’ attraction.
 The effective nuclear attraction for these outer
electrons is less, and so the outer electrons are less
tightly held.
 This makes it easier to remove some of these e-
when bonding to other atoms.

26. Ionization Energy
 This is the second important periodic trend.
 If an electron is pulled toward another atom with enough
energy to overcome the effective nuclear attraction of the
parent atom holding the electron in the cloud, it can leave the
atom completely.
 Once an electron has been removed from one atom and added
to another, both atoms are changed.
 These atoms have been “ionized”.
 The number of protons and electrons is no longer equal.
 The atom that lost an electron becomes POSITIVELY charged
(more P+ than e-) while the atom that gained an electron
becomes NEGATIVELY charged (more e- than P+ ).

27. Ionization Energy
 The energy required to remove an electron from an atom is
ionization energy.
 The larger the atom is, the easier its electrons are to remove.
 Ionization energy and atomic radius are inversely proportional.
 Ionization energy is always endothermic, that is energy is added
to the atom to remove the electron.

28. Ionization Energy
 Draw arrows on your help sheet like this:
DECREASES
DOWNWAR
D
TOP TO
BOTTOM
INCREASES LEFT TO RIGHT

29. Electron Affinity
 What does the word ‘affinity’ mean?
 Electron affinity is the energy change that occurs
when an atom gains an electron (also measured in
kJ).
 Where ionization energy is always endothermic,
electron affinity is usually exothermic, but not
always.

30. Electron Affinity
 Electron affinity is exothermic if there is an empty
or partially empty orbital for an electron to
occupy.
 If there are no empty spaces, a new orbital or PEL
must be created, making the process
endothermic.
 This is true for the alkaline earth metals and the
noble gases.

31. Electron Affinity
+
+

32. Metallic Character
 This is simple a relative measure of how easily
atoms lose or give up electrons.
 Your help sheet should look like this:

33. Electronegativity
 Electronegativity is a measure of an atom’s attraction for another
atom’s electrons.
 It is an arbitrary scale that ranges from 0 to 4.
 The units of electronegativity are Paulings.
 Generally, metals are electron givers and have low
electronegativities.
 Nonmetals are are electron takers and have high
electronegativities.
 What about the noble gases?

34. Electronegativity
0

35. Overall Reactivity
 This ties all the previous trends together in one
package.
 However, we must treat metals and nonmetals
separately.
 The most reactive metals are the largest since they
are the best electron givers.
 The most reactive nonmetals are the smallest ones,
the best electron takers.

36. Overall Reactivity
0

37. The Octet Rule
 The “goal” of most atoms (except H, Li and Be) is to have an
octet or group of 8 electrons in their valence energy level.
 They may accomplish this by either giving electrons away or
taking them.
 Metals generally give electrons, nonmetals take them from
other atoms.
 Atoms that have gained or lost electrons are called ions.

38. Ions
 When an atom gains an electron, it
becomes negatively charged (more
electrons than protons ) and is called an
anion.
 In the same way that nonmetal atoms can
gain electrons, metal atoms can lose
electrons.
 They become positively charged cations.

39. Ionic Radius
 Cations are always smaller than the original atom.
 The entire outer PEL is removed during ionization.
 Conversely, anions are always larger than the
original atom.
 Electrons are added to the outer PEL.

40. Cation Formation
11p+
Na atom
1 valence electron
Valence e-
lost in ion
formation
Effective nuclear
charge on remaining
electrons increases.
Remaining e- are
pulled in closer to
the nucleus. Ionic
size decreases.
Result: a smaller
sodium cation, Na+

41. Anion Formation
17p+
Chlorine
atom with 7
valence e-
One e- is added
to the outer
shell.
Effective nuclear charge is
reduced and the e- cloud
expands.
A chloride ion is
produced. It is
larger than the
original atom.

42. THANK YOU
CREATED BY :- LAKSHYA
SHARMA
lakshya2912@gmail.com