Characteristics of atoms
Key features of atoms
• All atoms are electrically neutral
• All atoms of the same element contain the same
number of protons and neutrons
• Number of protons = atomic number
• Most of the mass of an atom comes from Protons
and Neutrons
• Hence, the relative atomic mass is #Protons +
#Neutrons
Isotopes
• All atoms of an element have the same
number of protons, however, the number of
Neutrons can vary.
• Atoms with the same number of Protons, and
different number of Neutrons are called
isotopes
• Carbon 14 (draw examples on board)
Ions
• Atoms can gain or loose electrons, when they
do they are called ions
• Atoms which gain an electron become
negative i.e. Cl + e-  Cl• Atoms which loose an electron become
positive i.e. Na  Na+ + e• Draw examples on board
How are electrons arranged around
the nucleus?
• Rutherford proposed that electrons moved in
circular orbits, however physics says that
electrons moving in circular orbits should emit
electromagnetic radiation (such as light).
• As the radiation is emitted, the electrons should
lose energy and spiral into the nucleus, therefore
‘killing’ the atom.
• Rutherford's model also didn’t explain why, when
heated, elements only emitted light at certain
wave lengths.
Emission spectrum tests
Niels Bohr
• Suggested that small particles such as atoms did
not follow the laws physics displayed by large
objects
• Therefore, electrons circled the nucleus without
loosing energy
• Further, he proposed electrons could only move
in fixed orbits of certain energy levels
• Electrons with low energy moved in a small orbit
close to the nucleus, electrons with a larger
energy moved in a larger orbit further from the
nucleus
Jumping electrons
• Heating an element can cause an electron to
‘jump’ to a higher energy level (orbit), because
the electron absorbs some of this energy.
• Shortly afterward the electron returns back to
its initial state, releasing a fixed amount of
energy.
• Electrons can return to their energy level in a
number of different ways.
Ionisation energy
• So.. Electrons may be removed from their
‘shells’ by applying the atoms with energy.
• The further the negative electrons are from
the positively charged nucleus, the easier they
are to remove.
• The energy required to remove an electron is
called ionisation energy
An example
• Sodium (Atomic number 11, mass number 23)
has 11 protons, 12 neutrons and 11 electrons
• The electrons from Sodium can be removed,
and the ionisation energy can be measured.
• Look at the diagram on the next page
Shells
• After looking at ionisation energies of many
elements, scientists proposed that electrons
are grouped into ‘energy levels’ called shells.
• Electrons in the same shell:
- Are the same distance from the nucleus
- Have the same energy
• The various shells can hold different numbers
of electrons
Shells
•
•
•
•
The first shell can hold 2 electrons
The second, 8
The third, 18
The fourth 32
• 2n2 where n = number of shells
Electronic configuration
Electron dot diagrams
• Notice anything about Potassium? Remember
how many electrons the 3rd shell can hold?
Limitations of the Bohr model
• Why do electrons move in a circular motion,
and not an eliptical one?
• Why do the ‘shells’ have specific energies?
• Why does each shell have a maximum of 8
electrons, even though shells above 3 can hold
more than 8?
**Electrons behave like waves
• In 1926 Erwin Schrodinger proposed electrons
behave like waves around a nucleus
• Instead of circling in a specific orbit, they move in
waves around the nucleus, acting like a cloud of
negative charge.
• Schrodinger used his model (and the
accompanying equation) to predict the energy
levels of hydrogen, and larger atoms.
Quantum mechanics
• In this model, instead of occupying shells in an orbit,
electrons occupy regions of space called orbital's.
• In this model we will call these orbital's (regions of
space), ‘shells’, and will give them the shell numbers 1,
2, 3, 4 etc
• Within these shells, there are energy levels of similar
energy called subshells, these are labelled s, p, d, f
• Energies are as follows s < p < d < f
• Subshells are made up or orbitals (regions of space
where electrons move)
Electrons in subshells
Pauli Exclusion Principle
• Each orbital may hold a maximum of 2
electrons
Memorise This Chart!!!!
Draw it in your sleep!!!!
Examples
• Lets figure out the electronic configurations of a
few elements together 
• Na, Fe, Ne
• Step 1: Find out how many electrons in the
atom by looking at the atomic number
• Step 2: Draw the table from the previous page
• Step three: Fill the subshells until you run out of
electrons
Ground and Excited states
• Ground states – The lowest possible energy
state is called the ground state (these are
what we have listed so far)
• Excited states – When energy is applied to an
atom, an electron in the outer orbitals can
move to a higher energy level. i.e. Helium:
Ground state – 1s2
Excited state - 1s1, 2s1
Timeline
Essential questions for this chapter!
1, 4, 5, 6, 7, 9, 10, 12,
13 ,14, 17, 19, 22,
24, 29, 30