This document provides an overview of atomic chemistry, including: - Atomic structure, with shells containing orbitals and suborbitals that hold electrons. Electron configurations describe how electrons fill these orbitals. - The periodic table organizes elements based on atomic number and properties like radius, ionization energy, and electronegativity that change predictably across and down the table. - Bonding occurs through ionic bonds between elements with a large electronegativity difference, polar bonds between elements with a medium difference, and covalent bonds between elements with a small difference.

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 going across the period, and
decreases going down the group.

7. Across the period
● Atoms get smaller
● Atoms get more
electronegative
● Atoms’ first
ionization energy
increases
Down the group
● Atoms get bigger
● Atoms get less
electronegative
● Atoms’ first
ionization energy
decreases

8. 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.
The result is that polar molecules contain partial charges, called dipoles.
● 3. Ionic - the difference in electronegativity is large (>1.7) so one atom actually takes the electrons off the
other one. The result is that ionic molecules contain full charges, called ions.

9. 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: ion-ion (electrostatic) forces, hydrogen-bonds, dipole-dipole forces
and Van der Waal forces.
● The melting and boiling point of any molecule is defined solely by 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.

10. 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

11. 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 reaction. Unlike in chemical reactions, it occurs in the
nucleus, and will actually change the element into a whole new one! (for example C14 will turn into N)

12. 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.