Topic 2
Atomic Structure
A Brief History
• 1807: John Dalton proposed atomic theory.
– All matter was made up of small number of different kinds of
atoms.
– Atoms are indivisible and indestructible, but could combine in
whole numbers to form compounds.
• 1897: J. J. Thompson
– When electricity is passed through a very low pressure tube,
rays from the negative electrode (Cathode) could be detected
at the other end.
– Using magnets he bent the ray using magnets
– Found that they had charge (neg.) and mass in a fixed ration
The electron was discovered
A Brief History Continued
• 1909: Ernest Rutherford
– Fired alpha particles through gold foil only a few atoms
thick.
– He expected them to only deflect in small angles according
to Thomson’s model.
– Because some particles bounced back, he developed a
different model which has a dense, positively charged
nucleus as the core.
Brief History Con’t
• Ernest Rutherford
• Accelerated alpha α
particles (helium nuclei) at
thing gold (Au) foil
• Helium nuclei is the center
portion of an He molecule
without electrons.(e-)
• Analogy: Like shooting a
howitzer through a piece of
paper
• They did something odd.
• Summary
• alpha α particles are He
atoms without e• e- = electrons
• Au = gold
• He = Helium
• Drawing
Flash Animation
http://www.mhhe.com/physsci/chemistry/essent
ialchemistry/flash/ruther14.swf
Models
• Plum pudding
• Nuclear atom
• Negative ‘chocolate
chips’ were thought to
be surrounded by
positive ‘cookie dough’
• All mass around a small
positive center. Neg.
Electron cloud around it
IB Core Objective
• 2.1.1 State the position of protons,
neutrons, and electrons in the atom.
• State: Give a specific name, value or
other brief answer without
explanation or calculation.
2.1.1 State the position of protons, neutrons,
and electrons in the atom.
• Subatomic particles: Proton, neutron, and electron.
• Nucleus: The center portion of the atom.
– Contains protons and neutrons
• Electrons: Occupy shells around the nucleus.
• Orbital: Where an electron MIGHT be found
• If the nucleus is one meter across,
then the electrons would be about
10 kilometers across.
• Much of the atom is empty space.
IB Core Objective
• 2.1.2 State the relative mass and
relative charge of protons, electrons
and neutrons.
• State: Give a specific name, value or
other brief answer without
explanation or calculation. (Obj 1)
2.1.2 State the relative mass and relative charge
of protons, electrons and neutrons.
• Nucleus contains all the positive charge and
nearly all the mass (>99.9%).
• amu stands for atomic mass unit
Symbol
N0
P+
e-
Mass
1 amu
1 amu
5 * 10-4 amu
Charge
No charge
+1
-1
IB Core Objective
• 2.1.3 Define the terms mass
number (A), atomic number (Z) and
isotopes of an element.
• Define: Give the precise meaning of
a word, phrase or physical quantity.
(Obj 1)
2.1.3 Define the terms mass number (A), atomic
number (Z) and isotopes of an element.
• Atomic number (Z) is the number of protons
• Mass number (A) is the sum of the # of P+ and N0
(the nucleus)
• When writing these values, use the standard
notation format below:
2.1.3 Define the terms mass number (A), atomic
number (Z) and isotopes of an element.
Isotopes
• Many elements are composed of slightly
different types of atoms.
• These are called isotopes.
• Isotopes all have the same number of protons,
but differ in the number of neutrons.
IB Core Objective
• 2.1.4 Deduce the symbol for an
isotope given its mass number and
atomic number
• Deduce: Reach a conclusion from the
information given. (Obj 3)
2.1.4 Deduce the symbol for an isotope given its mass
number and atomic number
Using the Periodic Table
• Periodic table helps us find
information on different
elements.
• Atomic Mass: Tells you the
AVERAGE weight of the atom
Atomic
Number: Tells
you number of
protons
2.1.4 Deduce the symbol for an isotope given its
mass number and atomic number
The following notation should be used
AX
Z
Z= Number of protons
A=Atomic mass number (protons + neutrons)
2.1.4 Deduce the symbol for an isotope given its mass
number and atomic number
• Deduce the symbols using standard notation
for the following elements:
H
Cl
C
Fe
IB Core Objective
• 2.1.5 Calculate the number of protons,
neutrons and electrons in atoms and ions
from the mass number, atomic number
and charge.
• Calculate: Find a numerical answer
showing the relevant stages in the
working (unless instructed not to do so).
2.1.5 Calculate the number of protons, neutrons
and electrons in atoms and ions from the mass
number, atomic number and charge.
• Number of protons = atomic
number
• Number of electrons =
Number of protons
– If the atom is neutral (no
charge)
• Number of neutrons =
(Atomic mass) – (Number of
protons)
2.1.5 Calculate the number of protons, neutrons
and electrons in atoms and ions from the mass
number, atomic number and charge.
Calculating Neutrons
• The atomic weight MUST BE rounded.
– 1.3 is rounded down 1
– 1.6 is rounded up 2
– 13.6  14
– 35.45  35
– Now you can use this weight to find the number
of neutrons!
2.1.5 Calculate the number of protons, neutrons
and electrons in atoms and ions from the mass
number, atomic number and charge.
• What is the approximate number
of neutrons in iron?
• 30 (remember, the atomic mass is
an average based on the
abundance (%) of isotopes.)
IB Core Objective
• 2.1.6 Compare the properties of the
isotopes of an element.
• Compare: Give an account of
similarities and differences between
two (or more) items, referring to both
(all) of them throughout.(Obj 3)
2.1.6 Compare the properties of the isotopes of
an element.
• Carbon dating is done using the isotope C-14.
The 14 is the atomic weight.
• # of protons =
• # of electrons =
• # of neutrons =
• What is the difference from the carbon on the
periodic table?
2.1.6 Compare the properties of
the isotopes of an element.
2.1.6 Compare the properties of the isotopes of
an element.
• Elements may have multiple isotopes each
with a different abundance that all contribute
to the relative atomic mass of the element.
• Calculating Relative atomic mass:
–
37Cl
& 35Cl have relative abundances of 25% and
75% respectively, what is the relative atomic
mass?
IB Core Objective
• 2.1.7 Discuss the uses of radioisotopes
• Discuss: Give an account including,
where possible, a range of arguments for
and against the relative importance of
various factors, or comparisons of
alternative hypotheses. (Obj 3)
2.1.7 Discuss the uses of radioisotopes
• Radioisotopes are isotopes which are
radioactive.
• To be radioactive means the nucleus is
unstable, and will emit subatomic particles or
gamma rays (electromagnetic radiation).
2.1.7 Discuss the uses of radioisotopes
Homework
Go onto the uaschemistry.pbworks site.
In the 11th grade folder, go to the “Uses of
Radioisotopes” page
Research a radioisotope that we use, and comment
what it is and how we use it on this page.
If something has already been commented on, it cannot
be repeated, so better to start early!
You can expand on someone else’s comments.
These should be done by Tuesday, November 3rd
IB Core Objective
• 2.2.1 Describe and explain the operation of a
mass spectrometer.
• Describe: Give a detailed account. (Obj 2)
• Explain: Give a detailed account of causes,
reasons or mechanisms. (Obj 3)
2.2.1 Describe and explain the operation of a
mass spectrometer.
• Sequence for Mass
Spectroscopy
•
•
•
•
•
•
1) Vaporize
2) Ionize
3) Accelerate
4) Deflect
5) Detect
This needs to be done
in a vacuum to prevent
molecules colliding.
IB Core Objective
• 2.2.2 Describe how the mass
spectrometer may be used to
determine relative, atomic masses
using the 12C scale.
• Describe: Give a detailed account.
(Obj 2)
2.2.2 Describe how the mass spectrometer may
be used to determine relative, atomic masses
using the 12C scale.
Relative Atomic Mass (Ar)
The weighted mean mass of all naturally
occurring isotopes of that element
relative to one twelfth of carbon-12.
2.2.2 Describe how the mass spectrometer may be
used to determine relative, atomic masses using the 12C
scale.
• Deflection –
The accelerated ions are deflected
into the magnetic field.
The
amount of deflection is greater
when:
• the mass of the positive ion is less
• the charge on the positive ion is
greater
• the velocity of the positive ion is less
• the strength of the magnetic field is
greater
2.2.2 Describe how the mass spectrometer may be
used to determine relative, atomic masses using the 12C
scale.
• If all the ions are traveling at the same velocity and
carry the same charge, the amount of deflection in a
given magnetic field depends upon the mass of the
ion.
• For a given magnetic field, only ions with a particular
relative mass (m) to charge (z) ration – the m/z
value – are deflected sufficiently to reach the
detector.
2.2.2 Describe how the mass spectrometer may be
used to determine relative, atomic masses using the 12C
scale.
• Detection – ions that reach the detector cause
electrons to be released in an ion-current detector.
• The number of electrons released, hence the
current produced is proportional to the number of
ions striking the detector.
• The detector is linked to an amplifier and then to a
recorder: this converts the current into a peak
which is shown in the mass spectrum.
IB Core Objective
• 2.2.3 Calculate non-integer relative
atomic masses and abundance of
isotopes from given data.
• Calculate: Find a numerical answer
showing the relevant stages in the
working (unless instructed not to do so).
2.2.3 Calculate non-integer relative atomic
masses and abundance of isotopes from given
data.
• Determine the
z
relative atomic
mass from this
data.
28
29
30
• First, what
would your
estimate be?
z
2.2.3 Calculate non-integer relative atomic
masses and abundance of isotopes from given
data.
(28 x 92.21) + (29 x 4.79) + (30 x 3.09)
100
=28.1
2.2.3 Calculate non-integer relative atomic
masses and abundance of isotopes from given
data.
• A mass spec chart for a sample of neon shows that it
contains:
– 90.9% 20Ne
– 0.17% 21Ne
– 8.93% 22Ne
Calculate the relative atomic mass of neon
You must show all your working!
A = 20.18
r
IB Core Objective
• 2.3.1 Describe the electromagnetic
spectrum.
• Describe: Give a detailed account.
(Obj 2)
2.3.1 Describe the electromagnetic spectrum.
2.3.1 Describe the electromagnetic spectrum
The Chemical Fingerprint for Atoms
• Energy levels are discrete levels where
electrons can exist
http://jersey.uoregon.edu/
vlab/elements/Elements.h
tml
• When they are excited, elements will emit a
characteristic color of light.
• What can we do to excite them?
IB Core Objective
• 2.3.2 Distinguish between a continuous
spectrum and a line spectrum.
• Distinguish: Give the differences between
two or more different items.
2.3.2 Distinguish between a continuous
spectrum and a line spectrum.
IB Core Objective
• 2.3.3 Explain how the lines in the
emission spectrum of hydrogen are
related to electron energy levels.
• Explain: Give a detailed account of
causes, reasons or mechanisms. (Obj 3)
Chemical Finger Print
Higher energy level
(Farther from the
nucleus)
• When
electrons fall from
Energy levels (ring
Excited
State
ground
proximity)
notto
accurately
state, a represented
photon is emitted
Lower energy level
(Closer to the nucleus)
n =6
n =5
n =4
n =3
n =2
•Every atom has a
different number
therefore each atom has
its own fingerprint
Atom nucleus has been omitted
n=5
n=4
n=3
n=2
n=1
IB Core Objective
• 2.3.4 Deduce the electron arrangement
for atoms and ions up to Z = 20
• Deduce: Reach a conclusion from the
information given. (Obj 3)
2.3.4 Deduce the electron arrangement for
atoms and ions up to Z = 20
• Most stable energy levels are closest to the
nucleus.
• Electrons fill the most stable levels before
filling the higher ones.
• There is a maximum number each energy level
can hold. The first holds two electrons, the
second holds eight.
• The electrons in the valence shell determine
the physical and chemical properties of the
element.
Electron Arrangement
• The Bohr model (2,8,8,2)
• Li
Be
B
C
Ne
IB HL Objective
• 12.1.1 Explain how evidence from first
ionization energies across periods accounts for
the existence of the main energy levels and
sub-levels in an atom.
• Explain: Give a detailed account of causes,
reasons or mechanisms. (Obj 3)
12.1.1 Explain how evidence from first ionization
energies across periods accounts for the existence of
the main energy levels and sub-levels in an atom.
• You will be explaining ionization energies to
the class on Tuesday.
• You will all need to participate in helping
teach this to the other students.
• You can pick how you teach them (as a class,
each of you work with a small group, etc.)
12.1.1 Explain how evidence from first ionization
energies across periods accounts for the existence of
the main energy levels and sub-levels in an atom.
Ionization Energy
• Energy required to rip off one electron from
an element in its gaseous state.
X(g)  X+(g) + e-
12.1.1 Explain how evidence from first ionization
energies across periods accounts for the existence of
the main energy levels and sub-levels in an atom.
First Ionization Energy
The energy needed to remove the first electron
from a neutral atom.
Na(g) → Na+(g) + e-
12.1.1 Explain how evidence from first ionization
energies across periods accounts for the existence of
the main energy levels and sub-levels in an atom.
Second Ionization Energy
Hmmm, what do you think this might be?
The energy needed to remove the second
electron.
Na+ → Na2+(g) + e-
IB HL Objective
• 12.1.2 Explain how successive ionization
energy data is related to the electron
configuration of an atom.
• Explain: Give a detailed account of causes,
reasons or mechanisms. (Obj 3)
12.1.1 Explain how evidence from first ionization
energies across periods accounts for the existence of
the main energy levels and sub-levels in an atom.
• Look at Table 7.2 on page 271 of your book.
• Work with a partner(s) and explain the trends
in ionization energies that you see, and why
you see them
IB HL Objective
• 12.1.3 State the relative energies of s,p,d and f
orbitals in a single energy level.
• State: Give a specific name, value or other
brief answer without explanation or
calculation. (Obj 1)
12.1.3 State the relative energies of s,p,d and f
orbitals in a single energy level.
• Electrons are arranged around the nucleus in
specific energy levels and sub-levels.
• The different sub-levels differ in the shape of
electron distribution.
12.1.3 State the relative energies of s,p,d and f
orbitals in a single energy level.
S Orbitals
• Look at page 231 at figure 6.16.
• What is it showing?
• What is it’s shape?
12.1.3 State the relative energies of s,p,d and f
orbitals in a single energy level.
• All s orbitals are spherical and symmetric.
• 1s orbital is the lowest-energy orbital.