atomic chemistry

1. Atomic Chemistry
Studentgrinds.ie chemistry lesson 1

2. Atomic chemistry - overview
● Atomic structures
● Shells, orbitals, suborbitals, and electrons
● The periodic table
● Atomic radius, first-ionization energy and electronegativity
● Bonding
● Molecular forces, melting/boiling points and solubility
● Radioactivity

3. Structure of an atom
Shells:
These contain orbitals,
which contain
suborbitals, which
contain electrons. This
part of the atom has a
negative charge and
here (in the orbitals) is
where 2 atoms interact to
form a bond.
Nucleus:
This contains
the protons
and neutrons.
It holds all of
the atom’s
mass and
positive
charge.
● Every atom should have the same amount of protons, neutrons and electrons (unless it is
an ion or an isotope!)
● A normal atom will be neutral (it has no charge) because the positive charge in the nucleus
will exactly cancel out the negative charge in the shells.

4. Shells, orbitals and suborbitals
● Shells: this is the same as a quantum energy level. It is a level of energy
around the nucleus where orbitals can exist. In each shell, there are a certain
amount or orbitals.
● Orbitals: ‘a region with high probability of finding an electron’. There are four
types, based on their shape and how many suborbitals they have: s, p, d and
f orbitals.
● Suborbitals: This is where the electrons are. There’s a max of 2 electrons in
each suborbital. An s orbital has only 1 suborbital; a p has 3, a d has 5 and an
f has 7.

5. Electron configurations
● Because every element has a different number of
electrons, its electrons will be arranged slightly
differently around the atom.
● Step 1: how many electrons does the atom have? Just
look at the atomic number, and that’s your answer!
● Step 2: start filling up the orbitals (in the order of the
arrows on the diagram on the right) until you’ve used
up all the electrons. Remember: an s orbital has a max
of 2 electrons, a p has a max of 6, a d 10 and an f 14.
● Example: Sodium (Na) has 11 electrons. So when we
fill up the orbitals, in order of the arrows on the left, we
1 s
2 s p
3 s p d
4 s p d f
5 s p d f
get 1s2 2s2 2p6 3s1.
Energy level/ shell
Type of orbital
The amount of electrons in that
orbital
Don’t worry. This can seem quite confusing, but we’ll go through it in the live lesson

6. Trends in the periodic table
● Going across the table from left to right (called going across the ‘period’): we
add an extra proton. This increases the nuclear force (the atoms positive
attractive force).
● Going down the table (called going down the ‘group’): we add on an extra
shell, full of electrons.
● Atomic radius (size of the atom) : decreases going across period, and
increases going down the group.
● Electronegativity (how much the atom wants an extra electron) increases
going across the period, and decreases going down the group.
● First Ionization Energy (how hard it is to remove an electron from it) increases

7. Bonding (3 types)
● Electronegativity: a measure of how much an atom wants an extra electron
(not the actual definition!).
● The way any atoms will bond is defined by one thing and one thing only: the
difference in each atom’s electronegativity (which atom wants an extra
electron more, and by how much).
● 1. Covalent (non-polar) - the difference in electronegativity is very small (<0.4)
so both atoms share the electrons evenly.
● 2. Polar - the difference in electronegativity is medium (0.4-1.7) so they share
the electrons, but unevenly.
● 3. Ionic - the difference in electronegativity is large (>1.7) so one atom

8. Molecular forces (2 types)
● Intramolecular forces - hold atoms together within a molecule. In order of
strongest to weakest, these are: ionic bonds (the electrostatic force), polar
bonds, covalent bonds.
● Intermolecular forces - hold separate molecules together. In order of
strongest to weakest, these are: hydrogen-bonds, dipole-dipole forces and
Van der Waal forces.
● The melting and boiling point of any molecule is defined solely by, first, it’s
type of intramolecular forces; and, then, it’s type of intermolecular forces. The
stronger the forces are that hold something together, the more energy (heat)
will be required to break it apart, so it will have a higher melting/boiling point.

9. Solubility
● Solvent - dissolves a solute; solute - dissolved in a solvent. Together they
make up what’s called a ‘solution’.
● We use the term ‘like dissolves like’ - this actually refers to the type of
bonding. What it means is this:
● Any covalent solvent will dissolve anything else that’s covalent.
● Any polar solvent (i.e. water) will dissolve anything else that’s polar and
anything ionic.
● Ionic compounds are always solids, so they can’t really act as solvents

10. Radioactivity (3 types)
● Radioactivity: an unstable nucleus spontaneously decaying over time to
release radioactive material. There are 3 types of ‘radioactive material’:
● Alpha-particles (α): 2 protons and 2 neutrons, it’s basically the same as a
helium nucleus.
● Beta-particles (β): a high-energy electron, which actually comes from inside
one of the neutrons. This will turn that neutron into a proton!
● Gamma-rays (γ): very high energy light (‘light’ is also called electomagnetic
radiation)
● Nuclear reaction: most other reactions are called ‘chemical’ reactions, but
when an atom releases radioactive material from its nucleus, this is a nuclear

11. Also in Atomic Chemistry...
● The history of chemistry
● Definitions
Atomic chemistry is quite a big topic. Apart from all of the concepts above (which
will be explained during the live lesson), it also contains a list of scientists and
what they’re known for, and several definitions.
The reason that these are not included here is that they’re just a matter of learning
off chunks of text, and they’ll all be in your book; whereas this lesson focuses on
actually understanding the material and answering exam questions.