Atomic Structure
IB Chemistry Topic 2
The Atom
Originally thought to be the smallest form of
matter due to the fact that the atom can’t be
broken down into simpler components by chemical
reaction.
Dalton’s Atom
• All
matter is composed of tiny indestructible
particles called atoms
• Atoms
cannot be created or destroyed
• Atoms
of the same element are alike in every
• Atoms
of different elements are different
way
• Atoms
can combine together in small numbers to
form molecules (compounds)
Dalton’s Atom
• Dalton’s
Law of Constant Composition is seen
in the chemical formulas of compounds
• Each
atom is represented by its element symbol
• The
number of each type of atom is indicated by
a subscript written to the right of the symbol
• H2 O
N 2 O5
C6H12O6
Fe2(CO3)3
J.J. Thompson
• J.J.
Thomson (1856-1940) –atoms
contain negative particles (electrons)
• Proposed
the “plum pudding” model
of the atom - if a negative charge was
present in atoms, there must also be a
positive to balance it; negative charges
were distributed evenly on a positive
sphere (raisins suspended on pudding)
Gold Foil Experiment
• Ernest
Rutherford (1871-1937)
aimed positively charged
particles at a thin metal foil;
most particles went through
but some were deflected in
different directions
•
(a)If the plum pudding model was correct.
•
(b)Actual results
Structure of the Atom
• Protons + neutrons = nucleons  located in the nucleus
• Electrons are found in energy levels or shells
surrounding the nucleus
• Most of the atom is empty space
Structure of the Atom
Describing the atom
• Mass
Number (A) = number of protons + number
of neutrons in an element
A
Z
X
n+ / n-
Charge if it
is an ion,
zero if not
• Atomic Number (Z) = the number of protons
(also electrons in a neutral atom). The atomic
number is unchanging, it is the unique
identifier of an atom.
Label A and Z for the following:
Be
Ge
O
Te
Mo
Rb
Isotopes
• Atoms
of the same element (same atomic #) with different
masses due to a different number of neutrons .
• Examples:
 1H, 2H, 3H
 C-12, C-14
 35Cl, 37Cl
Deduce the symbol for an isotope given its
mass number and atomic number.
1. p = 6
n=7
2. p = 3
n=4
3. p = 16
n = 17
e = 18
Calculate the number of protons, neutrons and electrons in atoms
and ions from the mass number, atomic number and charge.
1. 17 Cl
35
e=?
p=?
n=?
2. Which subatomic particle occurs in the same amount in
both of the following species?
31
15
P
32
16
S
2-
Properties of Isotopes
• All
isotopes of an element have the same
chemical properties because chemical properties
are determined by the protons and electrons, not
the mass.
• Isotopes
have different physical properties.
 Ex: rate of diffusion, mass, density, melting
point, boiling point.
Radioisotopes
• Some
isotopes are radioactive (the nuclei
of these atoms break down spontaneously
and emit radiation).
Radioisotopes
• Uses:
• Nuclear power generation, sterilization of
surgical instruments in hospitals, crime
detection, finding cracks and stresses in
structural materials, food preservation…
• Add to your notes:
Iodine – 131, Cobalt – 60, Carbon – 14,
(brief description and uses)
• Shroud of Turin
• Pgs. 44-46
The electromagnetic spectrum
•
ℎ𝑐
𝐸 = ℎ𝜐 =
𝜆
Visible light only makes up a small part of the electromagnetic
spectrum.
Different types of spectra
A continuous spectrum is produced
when white light is passed through a
prism and shows all the frequencies.
Line spectra
• When white light passes through
gases certain absorption will occur
• Results in a line spectrum produce
with some colors of continuous
spectrum missing
• Different elements have different line
spectra
• Therefore elements can be identified
by their line spectra similar to
products with their barcodes
Transitions and EMS
Hydrogen spectrum:
The type of radiation given out by an atom is
dependent on where an electron falls to from its
excited state
Series
nf
ni
Region of EMS
Lyman
1
2,3,4,5…
UV
Balmer
2
3,4,5,6…
Visible and UV
Paschen
3
4,5,6,7…
IR
You don’t have to know
the series name, but you
do need to know which
region of the EMS goes
with which transitions.
Balmer:
Heisenberg’s uncertainty principle:
It is impossible to accurately know both the position
and the momentum of an electron. The more we know
about the position of an electron, the less we know about the
momentum, and vice versa.
Schrödinger’s equation:
A very complex math equation whose results are describes by
atomic orbitals. The results describe the probability density of
the space an electron can occupy.
An atomic orbital is a region in space where there is a high
probability of finding an electron.
Orbitals
• Energy
shells are divided into 4 levels:
 Principal quantum number, n, has integer values (1, 2,
3…) and can hold 2n2 electrons. This is the main energy
level.
 Sublevels, l: s, p, d, and f. Each of these levels can hold
2 electrons, but the number of orbitals, ml, in each
sublevel is different:
Sublevel
# of orbitals
Max e- number
s
1
2
p
3
6
d
5
10
f
7
14
Orbital Shapes
Orbitals
•
Orbitals in sublevel p are labeled px, py, and pz.
 **Labels for d and f do not need to be known.
•
Within each orbital, are two electrons that
have spins, ms, with values of +1/2 or -1/2.
•
Pauli exclusion principle: an orbital can hold
two electron and they must have opposite
spins.
•
Hund’s rule: electrons will fill each
degenerate orbital (orbitals of equal energy)
singly, before occupying them pairs.
Orbital Diagrams
When drawing orbital
diagrams, we use the arrowsand-boxes method.
For each sublevel, there is
one box per orbital.
For s: 1 box
For p: 3 boxes
For d: 5 boxes
For f: 7 boxes
Orbital diagrams
Aufbau principle: electrons
fill the lowest-energy orbital
that is available first.
Ex: Sulfur
16 electrons
Electron Configuration
• Full
electron configuration:
 Sulfur: 1s12s22p63s23p4
• Electron
configuration shorthand:
 Sulfur: [Ne]3s23p4
Remember this from CP? 