Download Atomic Theory

Atomic Theory
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2.1.1 State the position of protons, electrons and neutrons 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
2.1.5 Calculate the number of protons, neutrons and electrons in atoms and ions
from the mass number, atomic number and charge.
2.1.6 Compare the properties of the isotopes of an element
2.1.7 Discuss the uses of radioisotopes
History of the atom
• Democritus (400 BC) suggested that the material world was made
up of tiny, indivisible particles
• atomos, Greek for “uncuttable”
• Aristotle believed that all matter was made up of 4 elements,
combined in different proportions
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•
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Fire - Hot
Earth - Cool, heavy
Water - Wet
Air - Light
• The “atomic” view of matter faded for centuries, until early
scientists attempted to explain the properties of gases
Re-emergence of Atomic Theory
John Dalton postulated that:
All matter is composed of
extremely small, indivisible
particles called atoms
2. All atoms of a given
element are identical (same
properties); the atoms of
different elements are
different
1.
3. Atoms are neither created nor
destroyed in chemical reactions, only
rearranged
4. Compounds are formed when
atoms of more than one element
combine
• A given compound always has
the same relative number and
kind of atoms
Atoms are divisible!
• By the 1850s, scientists
began to realize that the
atom was made up of
subatomic particles
• Thought to be positive
and negative
• How would we know this
if we can’t see it or touch
it?
Cathode Rays and Electrons
• Mid-1800’s scientists began to study electrical discharge
through cathode-ray tubes. Ex: neon signs
• Partially evacuated tube in which a current passes through
• Forms a beam of electrons which move from cathode to
anode
• Electrons themselves can’t be seen, but certain materials
fluoresce (give off light) when energised
Oh there you are!
• JJ Thompson observed that when a magnetic
or electric field are placed near the electron
beam, they influence the direction of flow
• opposite charges attract each other, and
like charges repel.
• The beam is negatively charged so it was
repelled by the negative end of the magnet
• http://www.chem.uiuc.edu/clcwebsite/video/C
ath.mov
• Magnetic field forces the beam to bend
depending on orientation
• Thompson concluded that:
• Cathode rays consist of beams of particles
• The particles have a negative charge
• Thompson understood that all matter was inherently
neutral, so there must be a counter
• A positively charged particle, but where to put it
• It was suggested that the negative charges were balanced
by a positive umbrella-charge
• “Plum pudding model” “chocolate chip cookie model”
Rutherford and the Nucleus
• This theory was replaced
with another, more modern
one
• Ernest Rutherford (1910)
studied angles at which a
particles (nucleus of helium)
were scattered as they
passed through a thin gold
foil
• http://www.mhhe.com/p
hyssci/chemistry/essenti
alchemistry/flash/ruther1
4.swf
Rutherford expected …
• Rutherford believed that the mass and positive
charge was evenly distributed throughout the atom,
allowing the a particles to pass through unhindered
a particles
Rutherford explained …
• Atom is mostly empty space
• Small, dense, and positive at the center
• Alpha particles were deflected if they got close
enough
a particles
+
The modern atom is composed of two regions:
• Nucleus: Containing
protons and neutrons, it is the
bulk of the atom and has a
positive charge associated with
it
•
Electron cloud:
Responsible for the majority of
the volume of the atom, it is
here that the electrons can be
found orbiting the nucleus
(extranuclear)
Major Subatomic Particles
Name
Symbol
Charge Relative Mass Actual Mass (g)
(amu)
Electron
e-
-1
1/1840
9.11x10-28
Proton
p+
+1
1
1.67x10-24
Neutron
no
0
1
1.67x10-24
• Atoms are measured in picometers, 10-12 meters
• Hydrogen atom, 32 pm radius
• Nucleus tiny compared to atom
• If the atom were a stadium, the nucleus would be a marble
• Radius of the nucleus is on the order of 10-15 m
• Density within the atom is near 1014 g/cm3
Elemental Classification
• Atomic Number (Z) = number of protons (p+) in the
nucleus
• Determines the type of atom
• Li atoms always have 3 protons in the nucleus, Hg always 80
• Mass Number (A) = number of protons + neutrons
[Sum of p+ and nº]
• Electrons have a negligible contribution to overall mass
• In a neutral atom there is the same number of electrons
(e-) and protons (atomic number)
Nuclear Symbols
• Every element is given a corresponding symbol which is
composed of 1 or 2 letters (first letter upper case, second
lower), as well as the mass number and atomic number
mass number
A
elemental symbol
atomic number
Z
E
• Find the
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number of protons
number of neutrons
number of electrons
atomic number
mass number
19
9
F
80
35
Br
184
74
W
Ions
• Cation is a positively charged particle. Electrons
have been removed from the element to form
the + charge.
ex: Na has 11 e-, Na+ has 10 e-
• Anion is a negatively charged particle. Electrons
have been added to the atom to form the –
charge.
ex: F has 9 e-, F- has 10 e-
Isotopes
• Atoms of the same element can have different numbers
of neutrons and therefore have different mass numbers
• The atoms of the same element that differ in the
number of neutrons are called isotopes of that element
1
1
H
Hydrogen-1
2
1
H
Hydrogen-2
3
1
H
Hydrogen-3
• When naming, write the mass number after the name of the
element
How heavy is an atom of oxygen?
• There are different kinds of oxygen atoms (different
isotopes)
•
16O, 17O, 18O
• We are more concerned with average atomic masses,
rather than exact ones
• Based on abundance of each isotope found in nature
• We can’t use grams as the unit of measure because the
numbers would be too small
• Instead we use Atomic Mass Units (u)
• Standard u is 1/12 the mass of a carbon-12 atom
• Each isotope has its own atomic mass
Calculating Averages
– Average = (% as decimal) x (mass1) +
(% as decimal) x (mass2) +
(% as decimal) x (mass3) + …
– Problem:
– Silver has two naturally occurring isotopes, 107Ag with a mass
of 106.90509 u and abundance of 51.84 % ,and 109Ag with a
mass of 108.90476 u and abundance of 48.16 % What is the
average atomic mass?
– Average = (0.5184)(106.90509 u) + (0.4816)(108.90476 u)
= 107.87 amu
Average Atomic Masses
• If not told otherwise, the mass of the isotope is the mass
number in amu
• The average atomic masses are not whole numbers
because they are an average mass value
• Remember, the atomic masses are the decimal numbers
on the periodic table
Properties of Isotopes
• Chemical properties are primarily determined by the
number of electrons
• All isotopes has the same number of electrons, so they
have nearly identical chemical properties even though
they have different masses.
• Physical properties often depend on the mass of the
particle, so among isotopes they will have slightly
different physical properties such as density, rate of
diffusion, boiling point…
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The isotopes of an element with fewer neutrons will have:
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Lower masses
Lower densities
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faster rate of diffusion
lower melting and boiling points
More Practice Calculating Averages
• Calculate the atomic mass of copper if copper has two
isotopes
• 69.1% has a mass of 62.93 amu
• The rest (30.9%) has a mass of 64.93 amu
• Magnesium has three isotopes
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78.99% magnesium 24 with a mass of 23.9850 amu
10.00% magnesium 25 with a mass of 24.9858 amu
The rest magnesium 26 with a mass of 25.9826 amu
What is the atomic mass of magnesium?
Radioisotopes
• Isotopes of atoms that have had an extra
neutron attached to their nucleus.
• Carbon-14 radioactive decay is used to measures
the date of objects.
– After 5700 years the amount of 14C will be half its
original value.
• Iodine-125 or 131 is used to monitor the activity
of the thyroid gland (b/c the thyroid tends to
absorb iodine)
• Cobalt-60 produces gamma rays (intense
radioactivity) and is used in radiation treatment
of cancer.
• Note: gamma rays are the shortest wavelength
on the electromagnetic spectrum. They are the
most dangerous and difficult to shield from.
2.2 The Mass Spectrometer
Mass Spectrometer
• The mass spectrometer is an instrument used:
• To measure the relative masses of isotopes
• To find the relative abundance of the isotopes in a sample
of an element
When charged particles pass
through a magnetic field, the
particles are deflected by the
magnetic field, and the amount of
deflection depends upon the
mass/charge ratio of the charged
particle.
Mass Spectrometer – 5 Stages
• Once the sample of an element has been placed in the
mass spectrometer, it undergoes five stages.
• Vaporisation – the sample has to be in gaseous form.
If the sample is a solid or liquid, a heater is used to
vaporise some of the sample.
X (s)  X (g)
or X (l)  X (g)
Mass Spectrometer – 5 Stages
• Ionization – sample is bombarded by a
stream of high-energy electrons from an
electron gun, which ‘knock’ an electron from
an atom. This produces a positive ion:
X (g)  X + (g) + e– an electric field is used to accelerate
the positive ions towards the magnetic field. The
accelerated ions are focused and passed through a
slit: this produces a narrow beam of ions.
Acceleration
Mass Spectrometer – 5 Stages
• 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
Mass Spectrometer
• If all the ions are travelling 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.
Mass Spectrometer
• 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.
Atomic Structure – Mass
Spectrometer
•
 Isotopes
of boron
m/z value
Relative
abundance %
11
10
18.7
81.3
Ar of boron = (11 x 18.7) + (10 x 81.3)
(18.7 + 81.3)
= 205.7 + 813
100
= 1018.7 = 10.2
100
Mass Spectrometer – Questions
• 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 work!
Mass Spectrometer – Questions
– 90.9% 20Ne
– 0.17% 21Ne
– 8.93% 22Ne
(90.9 x 20u) + (0.17 x 21u) + (8.93 x 22u)
100
Ar= 20.18u
Mass Spectrometer – Questions
Calculate the relative
atomic mass of
lead
You must show all
your work!
52.3
23.6
22.6
1.5
204 206 207 208
m/e
Mass Spectrometer – Questions
–
–
–
–
1.5% 204Pb
23.6% 206Pb
22.6% 207Pb
52.3% 208Pb
(1.5 x 204) + (23.6 x 206) + (22.6 x 207)+(52.3 x 208)
100
306 + 4861.6 + 4678.2 + 10878.4
100
Ar= 207.24
=
20724.2
100
2.3 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 radiation.
Electromagnetic Radiation
• Most subatomic particles behave as
PARTICLES and obey the physics of
waves.
Electromagnetic Radiation
wavelength
Visible light
Amplitude
wavelength
Ultaviolet radiation
Node
Wavelengths and energy
• Understand that different wavelengths of
electromagnetic radiation have different
energies.
• Waves have a frequency
• c=νλ
– c=velocity of wave (2.998 x 108 m/s)
– ν=(nu) frequency of wave, units are “cycles per sec”
– λ=(lambda) wavelength
Electromagnetic Spectrum
In increasing energy, ROY G BIV
Electromagnetic Spectrum
Long wavelength --> small frequency
Short wavelength --> high frequency
increasing
frequency
increasing
wavelength
Bohr’s Model
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Why don’t the electrons fall into the nucleus?
Move like planets around the sun.
In circular orbits at different levels.
Amounts of energy separate one level from
another.
Bohr postulated that:
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Fixed energy related to the orbit
Electrons cannot exist between orbits
The higher the energy level, the further it is
away from the nucleus
An atom with maximum number of electrons
in the outermost orbital energy level is stable
(unreactive)
Think of Noble gases
Those who are not shocked when
they first come across quantum
theory cannot possibly have
understood it.
(Niels Bohr on Quantum Physics)
Atomic Line Emission
Spectra and Niels Bohr
Bohr’s greatest contribution to science
was in building a simple model of the
atom. It was based on an understanding
of the LINE EMISSION SPECTRA of
excited atoms.
•Problem is that the model only works for
Hydrogen
Niels Bohr
(1885-1962)
How did he develop his theory?
• He used mathematics to explain the visible spectrum of
hydrogen gas
• Lines are associated with the fall of an excited electron back
down to its ground state energy level.
• http://www.mhhe.com/physsci/chemistry/essentialchemis
try/flash/linesp16.swf
Spectrum of White Light
The line spectrum
• electricity passed through
a gaseous element emits
light at a certain
wavelength
• Can be seen when passed
through a prism
• Every gas has a unique
pattern (color)
Line Emission Spectra
of Excited Atoms
• Excited atoms emit light of only certain
wavelengths
• The wavelengths of emitted light depend on
the element.
Spectrum of
Excited Hydrogen Gas
Line Spectra of Other
Elements
Line spectrum
Continuous line spectrum
• Bohr also postulated that an atom would
not emit radiation while it was in one of
its stable states but rather only when it
made a transition between states.
• The frequency of the radiation emitted
would be equal to the difference in
energy between those states divided by
Planck's constant.
Ehigh-Elow= hν = hc/λ
h=6.63 × 10–34 J s = Planck’s constant
E= energy of the emitted light (photon)
ν = frequency of the photon of light
λ = is usually stated in nm, but for calculations use m.
• This results in a unique emission spectra for each
element, like a fingerprint.
• electron could "jump" from one allowed energy
state to another by absorbing/emitting photons of
radiant energy of certain specific frequencies.
• Energy must then be absorbed in order to "jump" to
another energy state, and similarly, energy must be
emitted to "jump" to a lower state.
• The frequency, ν, of this radiant energy corresponds
exactly to the energy difference between the two states.
• In order for the emitted energy to be seen as light
the wavelength of the energy must be in between
380 nm to 750 nm
Bohr’s Triumph
• His theory helped to explain periodic law (the
trends from the periodic table)
• Halogens (gp.17 or group VII) are so reactive
because it has one e- less than a full outer orbital
• Alkali metals (gp. 1 or group I) are also reactive
because they have only one e- in outer orbital
Drawback
• Bohr’s theory did not
explain or show the
shape or the path
traveled by the
electrons.
• His theory could only
explain hydrogen and
not the more
complex atoms
Energy level populations
• Electrons found per energy level of the atom.
• The first energy level holds 2 electrons
• The second energy level holds 8 electrons (2 in s and 6
in p)
• The third energy level holds 18 electrons (2 in s, 6 in p
and 10 in d) There is overlapping here, so when we do
the populations there will be some changes.
That is as far as this course requires us to go!
Examples for group 1
• Li
• Na
• K
2.1
2.8.1
2.8.8.1
The Quantum Mechanical Model
• Energy is quantized. It comes in chunks.
• A quanta is the amount of energy needed to
move from one energy level to another.
• Since the energy of an atom is never “in
between” there must be a quantum leap in
energy.
• Schrödinger derived an equation that described
the energy and position of the electrons in an
atom
Quantum or Wave Mechanics
Schrodinger applied idea of ebehaving as a wave to the problem
of electrons in atoms.
He developed the WAVE
EQUATION
Solution gives set of math
expressions called WAVE
E. Schrodinger
FUNCTIONS, 
1887-1961
Each describes an allowed energy
state of an e-
Heisenberg Uncertainty
Principle
W. Heisenberg
1901-1976
• The problem of defining nature of
electrons in atoms was solved by
W. Heisenberg.
• He observed that one cannot
simultaneously define the
position and momentum (= m•v)
of an electron.
• If we define the energy exactly of
an electron precisely we must
accept limitation that we do not
know exact position.
A good site:
http://www.chemguide.co.uk/basicorg/bo
nding/orbitals.html
Electron Configuration
HL only
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 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.
S orbitals
• 1 s orbital for
every energy level
1s
2s
3s
• Spherical shaped
• Each s orbital can hold 2 electrons
• Called the 1s, 2s, 3s, etc.. orbitals
P orbitals
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Start at the second energy level
3 different directions
3 different shapes
Each orbital can hold 2 electrons
The D sublevel contains 5 D orbitals
• The D sublevel starts in the 3rd energy level
• 5 different shapes (orbitals)
• Each orbital can hold 2 electrons
The F sublevel has 7 F orbitals
• The F sublevel starts in the fourth energy level
• The F sublevel has seven different shapes (orbitals)
• 2 electrons per orbital
Summary
Starts at
energy
level
Sublevel
# of shapes
(orbitals)
Max # of
electrons
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
Electron Configurations
• The way electrons are arranged in atoms.
• Aufbau principle- electrons enter the lowest energy
first.
• This causes difficulties because of the overlap of
orbitals of different energies.
• Pauli Exclusion Principle- at most 2 electrons per
orbital - different spins
• Hund’s Rule- When electrons occupy orbitals of equal
energy they don’t pair up until they have to .
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
6d
5d
4d
3d
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
5f
4f
3d
3p • Phosphorous, 15 e- to
place
2p • The first to electrons go
into the 1s orbital
• Notice the opposite spins
• only 13 more
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
5f
4f
3d
3p • The next electrons go into
the 2s orbital
2p
• only 11 more
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The next electrons go
into the 2p orbital
2p
• only 5 more
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The next electrons go
into the 3s orbital
2p
• only 3 more
5f
4f
Increasing energy
7s
6s
5s
4s
7p
6p
6d
5d
5p
4d
4p
3p •
3s
2s
1s
2p •
•
•
5f
4f
3d
The last three electrons
go into the 3p orbitals.
They each go into
separate shapes
3 unpaired electrons
1s22s22p63s23p3
Orbitals fill in order
• Lowest energy to higher energy.
• Adding electrons can change the energy of the
orbital.
• Half filled orbitals have a lower energy.
• Makes them more stable.
• Changes the filling order
•
•
•
•
•
•
•
Write these electron
configurations
Titanium - 22 electrons
1s22s22p63s23p64s23d2
Vanadium - 23 electrons
1s22s22p63s23p64s23d3
Chromium - 24 electrons
1s22s22p63s23p64s23d4 is expected
But this is wrong!!
Chromium is actually
•
•
•
•
•
1s22s22p63s23p64s13d5
Why?
This gives us two half filled orbitals.
Slightly lower in energy.
The same principal applies to copper.
•
•
•
•
•
Copper’s electron configuration
Copper has 29 electrons so we expect
1s22s22p63s23p64s23d9
But the actual configuration is
1s22s22p63s23p64s13d10
This gives one filled orbital and one half filled
orbital.
• Remember these exceptions
Electronic Structure – of transition metals
• With the transition metals it is the 4s electrons
that are lost first when they form ions:
– Titanium (Ti) - loss of 2 e-
1s2 2s2 2p6 3s2 3p6 4s2 3d2  1s2 2s2 2p6 3s2 3p6 3d2
Ti2+ ion
Ti atom

Chromium (Cr) - loss of 3 e-
1s2 2s2 2p6 3s2 3p6 4s1 3d5 
Cr atom
1s2 2s2 2p6 3s2 3p6 3d3
Cr3+ ion
Electronic Structure - Questions
• Copy and complete the following table:
Atomic
no.
Mass
no.
No. of
No. of
No. of
protons neutrons electrons
Mg
12
Al3+
27
S2Sc3+
Ni2+
21
1s2 2s2 2p6 3s2
10
16
16
45
30
Electronic
structure
26
Ionization Energy
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
Ionization Energy
• The amount of energy required to completely
remove an electron from a gaseous atom.
• An atom's 'desire' to grab another atom's
electrons.
• Removing one electron makes a +1 ion.
• The energy required is called the first ionization
energy.
X(g) + energy →X+ + e-
Ionization Energy
The second and third ionization energies can be
represented as follows:
• X+ (g) + energy X2+ (g) + e• X2+ (g) + energy X3+ (g) + e• More energy required to remove 2nd electron,
and still more energy required to remove 3rd
electron
Group trends
• Ionization energy decreases down the group.
Going from Be to Mg, IE decreases because:
– Mg outer electron is in the 3s sub-shell rather than
the 2s. This is higher in energy
– The 3s electron is further from the nucleus and
shielded by the inner electrons
– So the 3s electron is more easily removed
• A similar decrease occurs in every group in the
periodic table.
Notice any trends? Any surprises?
• General trend: Increasing I.E. as we go across a
period
• Look at the peak at Mg and the plateau between P and
S. Can you explain why?
Why is there a fall from Mg to Al?
• Al has configuration 1s2 2s2 2p6 3s2 3p1, its
outer electron is in a p sublevel
• Mg has electronic configuration
1s22s22p63s2.
• The p level is higher in energy and with Mg the
s sub level is full – this gives it a slight stability
advantage
Why is there a fall from P to S?
•
•
This can be explained in terms of electron pairing.
As the p sublevel fills up, electrons fill up the vacant
sub levels and are unpaired.
• This configuration is more energetically stable than S as
all the electrons are unpaired. It requires more energy to
pair up the electrons in S so it has a lower Ionization
energy.
• There is some repulsion between the paired electrons
which lessens their attraction to the nucleus.
• It becomes easier to remove!
Driving Force
• Full Energy Levels are very low energy.
• Noble Gases have full energy levels.
• Atoms behave in ways to achieve noble gas
configuration.
2nd Ionization Energy
• For elements that reach a filled or half filled
sublevel by removing 2 electrons 2nd IE is lower
than expected.
• Makes it easier to achieve a full outer shell
• True for s2
• Alkaline earth metals form +2 ions.
rd
3
IE
• Using the same logic s2p1 atoms have an low 3rd
IE.
• Atoms in the aluminum family form +3 ions.
•
nd
2
rd
3
IE and
IE are always
st
higher than 1 IE!!!