3.1 Periodic table
1. The periodic table consists of groups (vertical columns) and periods (horizontal rows).
Group: vertical row
Period: horizontal rows
2. Elements in the same period: the same number of shells occupied by electrons
Elements in the same group: the same number of valence electrons
3. Group number: the number of valence electron except group 0
Period number: the outer energy level that is occupied by electrons.
4.Categories of elements: dividing line separates metals and nonmetals
Metals: shiny, good electrical conductor
Non-metals: poor electrical conductor
Metalloids: having a mixture of metal and nonmetal properties
Example: Si has metallic luster but its conductivity of electricity changes with temperature
5. The periodic table is arranged into four blocks associated with the four sub-levels—s, p, d, and
f.s block: elements with outer electrons in s sub shell
p block: elements with outer electrons in p sub shell
d block: elements with electrons added to d sub shell
f block: elements with electrons added to f sub shell
6. Terms
Alkali metals: metals in group 1
Alkali earth metals: metals in group 2
Transition metals: metals that have an incomplete d block or that can form cations with an
incomplete d block
Halogens: elements in group 7
Nobel gases: elements in group 0
Lanthanoids (Lanthanides) : lanthanum and elements in the first row of f block (atomic number
increasing from 57 to 71)
Actinoids (Actinides) : actinium and elements in the second row of f block (atomic number
increasing from 89 to103)
3.2 Periodic trends
Periodicity: repeating pattern of (physical and chemical) properties
Nuclear charge: shows total positive charge of nucleus.
Value of nuclear charge= atomic number
Shielding effect: inner-shell electrons cancel part of the nuclear charge
Effective nuclear charge: the net positive charge of nucleus felt by the outer electron
Effective nuclear charge = Nuclear charge - the number of inner electrons
3.2.1 Periodic trends of physical properties
1. Electronegativity：a measure of attraction between an atom and its shared electron pairs with
other atoms in a covalent bond
Trend of electronegativity: increases across the period and decreases down the group
Electronegativity increases across the period because
 increasing nuclear charge
 similar shielding effect
 decreasing atomic radius
 So there is greater attraction for the electron pair in bond
Electronegativity decreases decreases down the group because
 bonding electrons are further away from nucleus
 similar effective nuclear charge: greater shielding effect despite greater nuclear charge
 So there is smaller attraction for the electron pair in bond
2. Ionization energy
1) First ionization energy: the energy required to remove one electron from each atom in one
mole of atoms of an element in the gaseous state.
Example: Na(g) ——> Na+(g) + e-
2) General trend of the first ionization energy: increases across the period and decreases down
the group
First ionization energy increases across the period because
•
increasing nuclear charge
•
similar shielding effect
•
decreasing atomic radius
So there is greater attraction between nucleus and outer electrons.
First ionization energy decreases down the group because
•
outer electrons are further away from nucleus
•
similar effective nuclear charge: greater shielding effect despite greater nuclear charge
So there is smaller attraction between nucleus and outer electrons
3) Two discontinuities
Compare first ionization energy of Be and B and explain
Be has greater first ionization energy because
•
Outer electron in Be is in 2s but is in 2p in B.
•
Electron in 2p is further away from nucleus/more shielded.
Compare first ionization energy of N and O and explain
N has greater first ionization energy because
•
Electrons singly fill 2p orbitals in O but are paried in one 2p orbitals in N.
•
Electron in 2p is further away from nucleus/more shielded.
4. Atomic radius
Bonding radius/covalent radius: half of the distance between nuclei of two atoms bonded to each
other
Non-bonding radius/ van der Waals’ radius: The shortest distance separating the two nuclei
during such collisions is twice the radii of the atoms.
Atomic radius increases down the group because
•
outer electrons in an energy level further away from the nucleus
•
so there is smaller attraction for the outer energy level of electrons.
Atomic radius decreases across the period because
•
nuclear charge increases
•
electrons are added to the same main energy level
•
so there is greater attraction for the outer energy level of electrons.
5. Ionic radius
1)
Positive ion is smaller than its parent atom.
Worked example: Compare the radius of a sodium cation and a sodium atom and explain
A sodium cation has smaller radius than a sodium atom because
•
A sodium atom has 11 electrons in three electron shells.
•
A sodium cation has 10 electron in two electron shells.
•
•
They have the same number of protons but sodium atom has one more electron and one more
shell occupied by electrons.
So the attraction for the outer energy level is smaller for sodium atom and radius of it is
bigger.
2) Negative ion is larger than its parent atom.
Worked example: Compare the radius of an oxide anion and an oxygen atom and explain
An oxide ion has bigger radius than an oxygen atom because
•
An oxygen atom has 8 electrons in two shells.
•
An oxide ion has 10 electron in two shells.
•
•
They have the same number of protons but the oxide ions has two more electrons than its
number of protons
So there is greater repulsion between electrons in oxide ions which makes attraction for outer
level smaller.
3) For isoelectric series species with greater nuclear charge have smaller radius.
• The ionic radii decrease from group 1 to 4 for positive ions.
• The ionic radii decrease from group 5 to 7 for negative ions.
Worked example: Compare the radius of Na+ and Mg2+
•
Both Na+ and Mg2+ have ten electrons in two energy levels.
•
Nuclear charge is greater in Mg so outer level is held more strongly.