#### Transcript 11th Grade IB Chemistry

```Unit 12: Atomic Structure
IB Chemistry
IB Topic 2 and 12
IB Topics 2 and 12
 2.1.1 State the position of protons, neutrons and
electrons in the atom.
 2.1.2 State the relative masses and relative
charges of protons, neutrons and electrons.
 2.1.3 Define the terms mass number (A), atomic
number (Z) and isotopes of an element.
 2.1.4 Deduce the symbol for an isotope given its
mass number and atomic number.
IB Topics 2 and 12
 2.1.5 Calculate the number of protons, neutrons
and electrons in atoms and ions from the mass
number, atomic number and charge.
HISTORY OF THE ATOM
460 BC
Democritus develops the idea of atoms
he pounded up materials in his pestle and
mortar until he had reduced them to smaller
and smaller particles which he called
ATOMA
(Greek for indivisible)
HISTORY OF THE ATOM
1808
John Dalton
suggested that all matter was made up of
tiny spheres that were able to bounce around
with perfect elasticity and called them
ATOMS
Dalton’s Atomic Theory (1808)
1.
Elements are composed of extremely small particles
called atoms. All atoms of a given element are identical.
The atoms of one element are different from the atoms of
all other elements.
2.
Compounds are composed of atoms of more than one
element. The relative number of atoms of each element in
a given compound is always the same.
3. Chemical reactions only involve the rearrangement of
atoms. Atoms are not created or destroyed in chemical
reactions.
HISTORY OF THE ATOM
1898
Joseph John Thompson
found that atoms could sometimes eject a far
smaller negative particle which he called an
ELECTRON
J.J. Thomson, measured mass/charge of e(1906 Nobel Prize in Physics)
Measured mass of e(1923 Nobel Prize in Physics)
e- charge = -1.60 x 10-19 C
Thomson’s charge/mass of e- = -1.76 x 108 C/g
e- mass = 9.10 x 10-28 g
HISTORY OF THE ATOM
1904
Thompson develops the idea that an atom was made up of
electrons scattered unevenly within an elastic sphere surrounded
by a soup of positive charge to balance the electron's charge
like plums surrounded by pudding.
PLUM PUDDING
MODEL
HISTORY OF THE ATOM
1910
Ernest Rutherford
oversaw Geiger and Marsden carrying out his
famous experiment.
they fired Helium nuclei at a piece of gold foil
which was only a few atoms thick.
they found that although most of them
passed through. About 1 in 10,000 were
deflected and, to their surprise, some helium
nuclei bounced straight back.
(1908 Nobel Prize in Chemistry)
 particle velocity ~ 1.4 x 107 m/s
(~5% speed of light)
1. atoms positive charge is concentrated in the nucleus
2. proton (p) has opposite (+) charge of electron
3. mass of p is 1840 x mass of e- (1.67 x 10-24 g)
Rutherford’s Model of
the Atom
atomic radius ~ 100 pm = 1 x 10-10 m
nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m
HISTORY OF THE ATOM
Rutherford’s new evidence allowed him to propose a more
detailed model with a central nucleus.
He suggested that the positive charge was all in a central
nucleus. With this holding the electrons in place by electrical
attraction
However, this was not the end of the story.
HISTORY OF THE ATOM
1913
Niels Bohr
studied under Rutherford at the Victoria
University in Manchester.
Bohr refined Rutherford's idea by adding
that the electrons were in orbits. Rather
like planets orbiting the sun. With each
orbit only able to contain a set number of
electrons.
Bohr’s Atom
electrons in orbits
nucleus
ATOMIC STRUCTURE
Particle
Charge
Mass
proton
+ ve charge
1
neutron
No charge
1
electron
-ve charge
nil
Subatomic Particles
Mass
(g)
Particle
Electron (e
-
)
9.1 x 10
Proton (p)
1.67 x 10
Neutron (n)
1.67 x 10
Charge
(Coulombs)
-28
-24
-24
-1.6 x 10
+1.6 x 10
0
-19
-19
Charge
(units)
-1
+1
0
mass p = mass n = 1840 x mass e-
Subatomic Particles
 For IB purposes we use:
 1 atomic mass unit (amu) as the mass of the
proton and neutron
 5x10-4 amu for the mass of the electron.
Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes are atoms of the same element (X) with different
numbers of neutrons in the nucleus
Mass Number
A
ZX
Atomic Number
1
1H
235
92
2
1H
U
Element Symbol
(D)
238
92
3
1H
U
(T)
An ion is an atom, or group of atoms, that has a net
positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation.
Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl
17 protons
17 electrons
Cl-
17 protons
18 electrons
A monatomic ion contains only one atom
Na+, Cl-, Ca2+, O2-, Al3+, N3-
A polyatomic ion contains more than one atom
OH-, CN-, NH4+, NO3-
Do You Understand Ions?
+
27
3
How many protons and electrons are in 13 Al ?
13 protons, 10 (13 – 3) electrons
2- ?
Se
How many protons and electrons are in 78
34
34 protons, 36 (34 + 2) electrons
SUMMARY
1. The Atomic Number of an atom = number of
protons in the nucleus.
2. The Atomic Mass of an atom = number of
Protons + Neutrons in the nucleus.
3.
The number of Protons = Number of Electrons.
4.
Electrons orbit the nucleus in shells.
5.
Each shell can only carry a set number of electrons.
Objectives
 2.1.6 Compare the properties of the isotopes of
an element.
 2.1.7 Discuss the uses of radioisotopes.
Isotopes
Atoms with the same number of protons, but
different numbers of neutrons.
Atoms of the same element (same atomic
number) with different mass numbers
Isotopes of chlorine
35Cl
17
chlorine - 35
37Cl
17
chlorine - 37
Learning Check
Naturally occurring carbon consists of three
isotopes, 12C, 13C, and 14C. State the number of
protons, neutrons, and electrons in each of these
carbon atoms.
12C
6
13C
6
14C
6
#p _______
#n _______
_______
_______
_______
_______
#e _______
_______
_______
Solution
12C
14C
13C
6
6
6
#p
6
6
6
#n
6
7
8
#e
6
6
6
Learning Check
An atom of zinc has a mass number of 65.
A.
Number of protons in the zinc atom
1) 30
2) 35
3) 65
B. Number of neutrons in the zinc atom
1) 30
2) 35
3) 65
C. What is the mass number of a zinc isotope
with 37 neutrons?
1) 37 2) 65
3) 67
Solution
An atom of zinc has a mass number of 65.
A.
Number of protons in the zinc atom
1) 30
B. Number of neutrons in the zinc atom
2) 35
C. What is the mass number of a zinc isotope
with 37 neutrons?
3) 67
Learning Check
Write the atomic symbols for atoms with the
following:
A. 8 p+, 8 n, 8 eB.
17p+, 20n, 17e-
C. 47p+, 60 n, 47 e-
___________
___________
___________
Solution
A. 8 p+, 8 n, 8 eB.
17p+, 20n, 17e-
16O
8
37Cl
17
C. 47p+, 60 n, 47 e-
107Ag
47
Atomic Mass on the Periodic
Table
Atomic Number
Symbol
Atomic Mass
11
Na
22.99
Atomic Mass
Atomic mass is the weighted
average mass of all the atomic masses
of the isotopes of that atom.
Example of an Average Atomic
Mass
natural chlorine.
35 x 75.5
= 26.425
100
37 x
24.5
100
=
9.065
35.49
Which isotope is it?
1.
A = 15, Z = 8
2.
A = 36, Z = 17
3.
A = 235, Z = 92
Which isotope is it?
1.
A = 15, Z = 8
2. A = 36, Z = 17
3. A = 235, Z = 92
 Radioactive isotopes of all elements can be
produced by exposing the natural element to a
flux of slow moving neutrons in a nuclear reactor.
 This makes the nucleus capture an extra neutron.
 The stability of a nucleus depends on the
balance between the number of protons and
neutrons.
 When a nucleus has too many or too few
neutrons it is radioactive and changes to a more
stable nucleus by giving out radiation.
 The radiation have different forms:
 Alpha particles: emitted by nuclei with too many
protons to be stable. They have 2 protons and 2
neutrons (the same as a helium nucleus).
 Beta particles: emitted by nuclei with too many
neutrons. They are electrons ejected from the
nucleus as a neutron decays.
 Gamma rays are a form of electromagnetic
 These radioisotopes have many uses:
 Generate energy in nuclear power stations
 Sterilize surgical instruments in hospitals
 Preserve food
 Fight crime
 Detect cracks in structural materials.
Carbon-14 Dating
 The radioactive decay of carbon-14 is used to date
carbon containing materials ex. fossils and
archaeological objects.

14C
is naturally present in the atmosphere and the
abundance in living things is constant – it is absorbed
from 14CO2 in the air.
 When organisms die the 14CO2 steadily decreases as
the radioactive isotope decays by releasing beta
particles.
 The half-life of the isotope is 5730 years so the amount
of 14C relative to 12C is compared to find how old the
item is.
Iodine-131as a Medical Tracer
 Iodine-131 emits beta and gamma radiation
and has a short half-life of 8 days. It is quickly
eliminated from the body.
 The activity of the thyroid gland can be
monitored when a person takes a drink that
contains iodine -131. This can also diagnose and
treat thyroid cancer.
 Iodine-125 emits beta radiation, has a half-life of
80 days and is used to treat prostate cancer by
implanting it in the gland.
cancerous cells. The treatment damages the
DNA in the cell making it impossible for the cell to
reproduce.

60Co
is an example of this. It emits gamma
 Radiotherapy is used to treat localized solid
tumors like skin, brain and breast; and blood
cancer like leukemia.
 Healthy cells can recover is dosage is controlled.
 What risks are associated with the use of
 How are these materials controlled?
Mass Spectrometer
 2.2.1 Describe and explain the operation of a
mass spectrometer.
 2.2.2 Describe how the mass spectrometer may
be used to determine relative atomic mass using
the 12C scale.
 2.2.3 Calculate the non-integer relative atomic
masses and abundance of isotopes from given
data.
Mass Spectrometer
Mass Spectrometer

Separates particles according to their masses
and records the abundance of each mass.

Process is as follows:
1.
Substance is vaporized and a vacuum is
maintained.
2.
Gaseous atoms or molecules are ionized to
have a positive charge by being bombarded
by electrons.
Mass Spectrometer
3.
Positive particles are accelerated by a large
potential difference between 2 electrodes.
4.
Fast moving particles pass through an
electromagnet which deflects them
according to their mass - the smaller particles
will be deflected more.
Mass Spectrometer
5.
Particles of a certain mass (adjustable) are
detected by the detector plate.
6.
The vacuum is there to prevent collision with
other particles.
Mass Spectrum
Calculating Relative Molecular Mass
mass of isotope1 x % abundance1 / 100
+
mass of isotope2 x % abundance2 / 100
=
relative molecular mass of element, Mr
Calculating Relative Molecular Mass
 A mass spectrum of chlorine shows there to be
25% 37Cl and 75% 35Cl. Calculate the relative
atomic mass of chlorine in this sample.
Calculating Relative Molecular Mass
 A mass spectrum of chlorine shows there to be
25% 37Cl and 75% 35Cl. Calculate the relative
atomic mass of chlorine in this sample.
 Solution:
(0.25 x 37) + (0.75 x 35) = 35.5
Calculating % abundance from
atomic mass
 Boron exists as
10.81
 If x = # of
10B
10B
and 11B. Its relative atomic mass is
atoms then 100 – x = # of 11B atoms
 Total mass = 10x + (100-x)11
= 10x + 1100 – 11x
= 1100 – x
 Average mass = total mass / number of atoms
= (1100-x) / 100
10.81 = 1100 – x
x = 1100 – 1081 = 19.00
 The abundances are
10B
= 19.00% and 11B = 81.00%
Electron Arrangement
 2.3.1 Describe the electromagnetic spectrum.
 2.3.2 Distinguish between a continuous spectrum
and a line spectrum.
 2.3.3 Explain how the lines in the emission
spectrum of hydrogen are related to electron
energy levels.
 2.3.4 Deduce the electron arrangement for
atoms and ions up to Z = 20.
Electromagnetic Spectrum
 Light is a general word used to describe all
wavelengths of electromagnetic radiation - from
short gamma rays to long radio waves.
 White light is made up of all the visible colors of
light.
 Electromagnetic radiation is commonly called
light waves.
 All light waves have the same speed (c) but
have different wavelengths (λ) and therefore
different frequencies (f).
 Frequency and wavelength are inversely
proportional since fewer long waves can pass
the same point in space in a certain period of
time.
 They are related by the equation c = fλ
 c = fλ
 Where c is the speed of light
 f is frequency measured in Herz (Hz)
  is wavelength measured in meters (m)
 c = 3 x 108 ms-1
Continuous Spectrum
 If white light is passed through a prism we see the
different colors that make up the light.
 The different wavelengths blend into each other
so this is called a continuous spectrum.
Evidence for orbitals
 The best evidence for the existence of energy
levels in an atom comes from studying emission
spectra of elements.
 When an element is excited it will often emit a
light of a characteristic color ex. Red light in
neon signs.
Colors
 Gases can be excited by passing electricity
through them at low pressure.
 Some metals will show a distinctive color when
heated in the bunsen flame.
 Flame tests are commonly used to identify alkali
metals.
Flame Colors
 Lithium - red
 Sodium - yellow
 Potassium - lilac
 Copper - green
 Barium - yellow-green
 Calcium - brick red
Line Spectrum
 If the light that is emitted is looked at through a
prism or diffraction grating, different lines of color
can be seen.
 These lines make up the line spectrum.
 Each element has its own specific line spectrum.
Emission vs Absorption Spectra
 An emission spectrum shows lines of color against
a black background. The lines of color are the
wavelengths of light that are produced.
 An absorption spectrum shows lines of black on
a colored background. The lines of black are the
wavelengths where light is absorbed.
Planck’s Equation
 When an atom is excited its electrons
gain energy and they move to a higher
energy level.
 When returning to the lower energy level
they lose this energy as light which we
see as colors.
 Only certain energy levels can be
reached so only certain values of energy
can be lost, therefore only certain
frequencies of light can be emitted.
Planck’s Equation
 The color of the light is measured by
frequency and is related to energy by the
equation:
 ∆E = hf
 where h is Planck’s constant, E is energy and
f is frequency (1Hz = 1s-1).
 Planck’s constant = 6.62 x 10-34 m2 kg / s
Atomic Emission Spectrum of
Hydrogen
 This is the simplest emission spectrum.
 Hydrogen only has 1 electron so there is no
repulsion from other electrons.
 The spectrum consists of bright lines.
 The lines are in series named after the scientists
who discovered them.
Hydrogen emission spectrum
Hydrogen Emission Series
 Each series has a similar structure of lines that
converge (get closer together) at higher
frequencies / higher energies.
 Each series corresponds to transitions in which
the electron falls to a particular energy level.
 The largest energy transition is from n = 2 to n = 1
and the lines are in the U.V. region (high energy)
 All transitions to the n = 2 level are in the visible
region.
 Transitions to n = 3 are in the infrared region.
Hydrogen Emission Spectrum
 Each series ends in a brief continuum at the high
frequency end where the lines become too
close together to be separated.
 When an electron is at n = ∞, it is no longer in the
atom and the atom has been ionized (positive
ion formed).
 In the case of the Lyman series this is the
ionization energy of hydrogen.
Electron Configuration
 12.1.1 Explain how evidence from first ionization
energies across periods accounts for the
existence of main energy levels and sub-levels in
atoms.
 12.1.2 Explain how successive ionization energy
data is related to the electron configuration of
an atom.
Trends in Ionization Energies
 Ionization energy is the energy required to
remove an electron from an atom to form a
positive ion. Units: kJmol-1.
 Ex. Al(g)  Al+(g) + e-
 The second ionization energy refers to
removing another electron from a 1+ ion:
 Al+(g)  Al2+(g) + e Note that all of these changes are measured
in the gas phase.
Trends in Ionization Energies
 The successive ionization energies of Aluminum
are shown below. Note the jump from the 3rd
electron to the 4th and again from the 12th
electron to the 13th.
Trends in Ionization Energies
 There is an overall increase in ionization energy
as more electrons are removed.
 The first ionization involves separating an
electron from a neutral atom.
 The second ionization involves removing an
electron from a positively charged ion.
 As more electrons are removed the remaining
electrons are more strongly attracted to the
increasingly positively charged ion – so more
energy is required.
Trends in Ionization Energies
 The large increases in successive ionization
energies indicate that the attractive force
holding the electrons is becoming stronger.
 This evidence led to the theory of different
energy levels within an atom.
 Electrons must be removed from outer energy
levels first.
 Then electrons in energy levels closer to the
nucleus require are removed but more energy is
required.
Trends in Ionization Energies
 Atomic theory explains that an atom has
discrete (separate) energy levels that hold a
certain number of electrons.
 As electrons are removed the sudden increases
in ionization energies indicate an electron has
been removed from an energy level nearer the
nucleus.
 If you look closely at ionization data you will see
smaller increases within one energy level
indicating the sub-levels that exist in atoms.
Energy Levels
 12.1.3 State the relative energies of s, p, d and f
orbitals in a single energy level.
 12.1.4 State the maximum number of orbitals in a
given energy level.
 12.1.5 Draw the shape of an s orbital and the
shapes of the px, py and pz orbitals.
 12.1.6 Apply the Aufbau principle, Hund’s rule
and the Pauli exclusion principle to write electron
configurations for atoms and ions up to Z = 54.
Energy Levels
 Electron configuration is a way to show where
each electron in an atom exists.
 We think of electrons as being in energy levels.
 The lowest energy level is closest to the nucleus.
 As we move further from the nucleus the energy
increases.
Energy Levels
 The main energy level is assigned a number with level
1 being closest to the nucleus. It is labeled n=1.
 Each main energy level contains 1 or more sub-level
which are labeled with the main energy number
followed by the type of orbitals it contains.
 We use the letters s, p, d and f to denote each type
of orbital.
 Ex. Energy level n = 2 contains sub-levels 2s and 2p.
Orbitals
 An orbital is a region in space around the
nucleus that an electron is likely to be found.
 It is only a probability given by electron density -
we cannot pinpoint the location of an electron
because it is always moving very quickly.
 Each orbital has a different 3-D shape and
capacity for electrons.
Heisenberg’s Uncertainty
Principle
 It is possible to think of light waves as discrete
particles – called photons.
 In the same way electrons (or any object) can
be thought of as behaving like waves.
 In reality the behavior is somewhere between
these two extremes.
 Heisenberg said that it is impossible to make an
exact measurement of both the position and
momentum of any object at the same time.
Heisenberg’s Uncertainty
Principle
 Locating an electron in an atom is an example
of this principle.
 The act of trying to find the electron using
radiation gives the electron a random “kick”
sending it off in a random direction.
 In other words, the more you know about an
electrons location the less you know about its
speed and vice versa.
Orbitals
 The s orbital has a spherical shape and each s
orbital can hold 2 electrons.
 When an electron is in an energy level it spins in
a particular direction.
 When there are two electrons in an orbital
together they must spin in opposite directions.
 This is a result of the PAULI EXCLUSION PRINCIPLE.
Orbitals
 p orbitals have 3 sub-levels and each can hold 2
electrons.
 This gives a total capacity of 6 electrons.
 The p orbitals have a dumb-bell shape in 3
dimensions.
Orbital Shapes
Electronic Structure
 IB describes how many electrons are in each
period using “electronic structure”.
 They simply write the number of electrons in
each period separated by a comma.
 Ex. H is 1, He is 2 and Li is 2,1
 Write the electronic structure for Be, Al, Kr
 Using the periodic table to count up electrons
Electron Configuration
 We use the orbitals to show how many electrons
an atom or ion has and where they are located.
 It is similar to a ticket for a sports event where
section, row and seat number are given for each
person.
How to write electron configuration
 The first part of the electron configuration
denotes which period (row) of the periodic
 The second part tells us the orbital, s,p, d or f.
 The third part is a superscript number that
tells us how many electrons are in the orbital.
 For example:
 1s1 is H
1s22s1 is Li
Learning Check

Write the electron configurations for the
following:
1.
He
2.
C
3.
F
Learning Check
 Write the electron configurations for the
following:
1. He 1s2
2. C 1s22s22p2
3. F 1s22s22p5
Rules for Electron Configuration
 The AUFBAU PRINCIPLE states that orbitals fill up
with electrons from the lowest energy first.
 HUND’S RULE states that the lowest energy is
achieved when the maximum number of
electrons have the same spin direction within an
energy level.
 Ex.    rather than   . . .
 IB expects you to use arrows in boxes to
represent electrons – PIC!
Order of Filling Orbitals
Order of Filling Orbitals
Electron Configuration and the
Periodic Table
 The periodic table is arranged according to
these sub-levels and the parts are referred as
blocks.
 Elements whose outer electrons occupy an s
sub-level are in the s block.
 Elements with outer electrons in the p orbitals
make up the p block and so on for d and f
blocks.
Periodic Table
 The periodic table can be used to write electron
configurations:
Ionization Energies… Again!
 There is a general increase across a period when
looking at first ionization energies.
 As the number of protons in the nucleus
increases it become more difficult to remove an
electron.
 There is then a decrease down to the next
period.
 Within the period there are small increases at the
change between s and p sub-levels and halfway through the p sub-level.
Ion Configurations
 Ions have different numbers of electrons from
their uncharged atoms so the electron
configuration must change too.
 A positive ion has less electrons so the electron
configuration will be shorter.
 A negative ion has more electrons so the
electron configuration will be longer.
Learning Check

Write the configurations of these ions:
1.
Na+
2.
Al3+
3.
Cl-
4.
O2-
Learning Check

Write the configurations of these ions:
1.
Na+
1s22s22p6
2.
Al3+
1s22s22p6
3.
Cl1s22s22p63s23p6
4.
O2-
1s22s22p6
```