IB CHEMISTRY:
TOPIC 3 PERIODICITY NOTES
3.1 Periodic Table
Periodic Table Arrangement
• The periodic table is arranged into 4 blocks associated with their last orbital: s, p,
d and f
• Elements in the periodic table are arranged by increasing atomic number
• Group: A vertical column of elements
o Group number is the same as the number of valence electrons
• Period: A horizontal row of elements
o Period number is the same number of shells in the atom
o All shells except the outer shell will be full
• There is a division between metals and non-metals. Metals are on the left and non-
metals are on the right
• The farther sub-level occupied is found by the location of the element (which block it is
in):
• The location of the following families of the periodic table must be known: alkali metals,
alkali earth metals, halogens, noble gases, transition metals, lanthanides and actinoids
3.2 Periodic trends
Definitions
Atomic Radius – The distance from the nucleus to the outermost electron
, Ionic Radius – The distance from the nucleus to the outermost electrons in an ion
First ionization energy – The energy required to remove one mole of electrons from one
+
mole of gaseous atoms. This is shown by: 𝑋(𝑔) → 𝑋(𝑔) + 𝑒−
Electron affinity – The energy released when one mole of an electron is added to one mole
−
gaseous atoms. This is shown by: 𝑋(𝑔) + 𝑒 − → 𝑋(𝑔)
Electronegativity – A measure of the attraction an atom has for a shared pair of electrons in a
covalent bond
Trends in the Periodic Table
• Periodicity refers to repeating trends or patterns of physical and chemical properties in
elements
Atomic Radius
• Atomic radius increases down a group as the number of
electron shells increases
• Atomic radius decreases across a period
o This is because electrons are added to the same main
energy level (n=3) the nuclear charge also increases
o The attraction between the nucleus and the outer
electrons increases resulting in a smaller radius
Ionic Radius
• Ionic radius of positive ions decrease across a period as the number of protons in
the nucleus increases but the number of electrons remain the same
• Ionic radius of negative ions decrease across a period as the number of protons in
the nucleus increases but the number of electrons remains the same
Ionization energy
• Ionization energy increases across a period
• The increase in nuclear charge across a period causes an increase in the attraction
between the outer electrons and the nucleus makes the electrons more difficult to
remove
• Ionization energy decreases down a group
• The electron being removed is from the energy level furthest
from the nucleus so it gets easier to remove valence electrons as
atomic radius increases down a group
o Valence Electrons: The outermost electrons of an atom
Electron affinity
• Generally, metals have a low EA and non-metals have a
higher EA
• The greater the distance between the nucleus and the outer energy level, the weaker the
electrostatic attraction and the less energy is released when an electron is added to the
atom
TOPIC 3 PERIODICITY NOTES
3.1 Periodic Table
Periodic Table Arrangement
• The periodic table is arranged into 4 blocks associated with their last orbital: s, p,
d and f
• Elements in the periodic table are arranged by increasing atomic number
• Group: A vertical column of elements
o Group number is the same as the number of valence electrons
• Period: A horizontal row of elements
o Period number is the same number of shells in the atom
o All shells except the outer shell will be full
• There is a division between metals and non-metals. Metals are on the left and non-
metals are on the right
• The farther sub-level occupied is found by the location of the element (which block it is
in):
• The location of the following families of the periodic table must be known: alkali metals,
alkali earth metals, halogens, noble gases, transition metals, lanthanides and actinoids
3.2 Periodic trends
Definitions
Atomic Radius – The distance from the nucleus to the outermost electron
, Ionic Radius – The distance from the nucleus to the outermost electrons in an ion
First ionization energy – The energy required to remove one mole of electrons from one
+
mole of gaseous atoms. This is shown by: 𝑋(𝑔) → 𝑋(𝑔) + 𝑒−
Electron affinity – The energy released when one mole of an electron is added to one mole
−
gaseous atoms. This is shown by: 𝑋(𝑔) + 𝑒 − → 𝑋(𝑔)
Electronegativity – A measure of the attraction an atom has for a shared pair of electrons in a
covalent bond
Trends in the Periodic Table
• Periodicity refers to repeating trends or patterns of physical and chemical properties in
elements
Atomic Radius
• Atomic radius increases down a group as the number of
electron shells increases
• Atomic radius decreases across a period
o This is because electrons are added to the same main
energy level (n=3) the nuclear charge also increases
o The attraction between the nucleus and the outer
electrons increases resulting in a smaller radius
Ionic Radius
• Ionic radius of positive ions decrease across a period as the number of protons in
the nucleus increases but the number of electrons remain the same
• Ionic radius of negative ions decrease across a period as the number of protons in
the nucleus increases but the number of electrons remains the same
Ionization energy
• Ionization energy increases across a period
• The increase in nuclear charge across a period causes an increase in the attraction
between the outer electrons and the nucleus makes the electrons more difficult to
remove
• Ionization energy decreases down a group
• The electron being removed is from the energy level furthest
from the nucleus so it gets easier to remove valence electrons as
atomic radius increases down a group
o Valence Electrons: The outermost electrons of an atom
Electron affinity
• Generally, metals have a low EA and non-metals have a
higher EA
• The greater the distance between the nucleus and the outer energy level, the weaker the
electrostatic attraction and the less energy is released when an electron is added to the
atom
