Download 2. Atomic structure

2. Atomic structure
“A wrong theory is always so much better than no theory
at all.” William Lawrence Bragg (1890-1971)
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Richard Feynman (1918-1988):
“If all of scientific knowledge were to be
destroyed, and only one sentence passed to
the next generation....I believe it is that all
things are made of atoms.”
Daltons atomic theory
1. All matter consists of very small
particles called atoms.
2. An element consists of atoms of one
type only.
3. Compounds consist of atoms of more
than one element and are formed by
combining atoms in whole-number ratios
(the law of multiple proportions).
4. In a chemical reaction atoms are not
created or destroyed.
Thomsons “plum-pudding” model of the atom
Rutheford´s gold foil experiment
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This experiment tested Rutherfords “plum-pudding
model” and led to the conclusion that the atom is
mainly empty space.
http://www.youtube.com/watch?v=wzALbzTdnc8&feature=results_main&playnex
t=1&list=PLDF89CB07DD4E004E
The atom
In a neutral atom, the number of electrons equals the number of protons.
Sub-atomic particle
Charge
Mass (kg)
Mass/amu
proton
+1
1,67 ▪ 10-27 kg
1
neutron
0
1,67 ▪ 10-27 kg
1
electron
-1
9,1 ▪ 10-31 kg
0,0005
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Mass number (A) = number of protons plus neutrons in
the nucleus of an atom
Atomic number (Z): the number of protons in the
nucleus of an atom.
Isotopes
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Atoms of the same element always contain the same
number of protons (= same atomic number Z), but if
they contain a different number of neutrons (=
different mass number A) they are called isotopes.
Most elements occur in nature as a mixture of
isotopes.
For most of the light elements, the numbers of
protons and neutrons in the nucleus are nearly equal.
Radioisotopes
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Are used in:
- Nuclear medicine for diagnostics, treatment and
research (iodine-131, iodine-125, cobalt-60)
- Tracers in biochemical and pharmaceutical
research
- Chemical clocks in geological and archaeological
dating (carbon-14)
https://www.youtube.com/watch?v=udkQwW6aLik
Half-life, t1/2
Radioactive decay
• Many isotopes of elements are radioactive, because their nuclei
break down spontaneously and emit radiation:
• alpha, α:
- two protons and two neutrons
- can not penetrate the skin
• beta, β:
- Stream of electrons that comes from the
nucleus, NOT from the electron shells
• gamma, γ: - High energy electromagnetic radiation
- Extremely penetrating
- Blocked efficiently by very dense materials: Pb
Relative atomic mass
• The mass of an atom depends on the number of protons
and neutrons in the nucleus (the mass of electrons is so
small that it can be ignored in chemistry).
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Since the mass of a single atom is very small (on
average 10-23 g), it is not possible to weight single atoms
or molecules.
The relative atomic mass of an element is the weighted
mean mass of all the naturally occuring isotopes of that
element relative to the mass of carbon-12.
The unified atomic mass unit
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One-twelfth of the mass of a carbon-12 atom in its
ground state.
1 amu or 1 u = 1.6605402 x 10-27 kg
The relative atomic mass, Ar
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The ratio of the average mass of the atom to the
unified atomic mass unit.
The mass spectrometer
• The relative atomic mass of an element can be
determined using a mass spectrometer:
1) the sample is vaporized
2) ionized to positive ions: M (g) → M+ (g) + e3) the ions are accelerated in an electric field
4) the ions are deflected by an external magnetic
field
5) the ions are recorded on a detector (relative
amounts of the ions, mass to charge ratio)
https://www.youtube.com/watch?v=mBT73Pesiog
2.2 Electron configuration
Metals often give different characteristic colours when heated
strongly, i.e. given energy.
e.g.
Cu:
Ba:
Na:
Li:
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Each element has its own characteristic colour which can be used
to identify the element.
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The electromagnetic spectrum (EMS)
• Visible light is one type of electromagnetic radiation.
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The peak to peak distance is called the radiation's
wavelength, λ.
• The number of waves which pass a point in one
second is the frequency of the radiation, v.
• Wavelength is related to the frequency of the
radiation:
c = vλ
c = speed of light ≈ 3,00 · 108 m s-1
v = frequency of the radiation (Hz)
λ = wavelength (m)
Continuous spectrum
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When white light (sunlight) is passed through a prism, a
continuous spectrum of all colours can be obtained.
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A continuous spectrum contains light of all wavelengths
and so appears as a continuous series of colours.
Line spectra
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When white light is passed through hydrogen gas,
some of the light is absorbed.
An absorption spectrum is produced where some
colours are missing ( those that are absorbed by
hydrogen).
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A corresponding emission line spectrum can be
produced. The lines correspond to the light of
particular wavelengths given off (=emitted) by the
element.
http://www.youtube.com/watch?v=nM5Kg7RUoTE
The hydrogen spectrum
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How can a hydrogen atom absorb and emit energy?
• Niels Bohr proposed a theoretical explanation:
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electrons travel in orbits around the nucleus
The energy of each orbit is fixed (=quantized)
http://www.youtube.com/watch?v=nM5Kg7RUoTE
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Light can be described as a stream of photons or tiny
“packets” (= quanta) of light energy.
The energy of electromagnetic radiation is expressed:
E = hv
E = energy of a photon (J)
h = Planck's constant ≈ 6,626 · 10-34 J s
v = frequency of the radiation (Hz)
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The smaller the wavelength (= the higher the
frequency), the greater the energy of the photon.
Excitation
• When an atom absorbs energy, an electron moves into an
orbit of higher energy further from the nucleus. An unstable
excited state is produced.
• The electron soon falls back to a lower level and gives out (=
emits) the energy in the form of electromagnetic radiation
with a specific frequency, v.
• The lowest energy level of the electron is called the ground
state.
∆Eelectron = Ef- Ei = hv
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The line spectrum of hydrogen provides evidence for
the existence of electrons in discrete energy levels,
which get closer together (=converge) at higher
energies.
Electron arrangement
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In the Bohr model of the atom the energy levels are
drawn as shells.
Each shell or orbit can be occupied by a certain number
of electrons.
Electron configuration
• The Bohr´s model of an atom is a simplification
that only explains the spectral lines of
hydrogen.
• Limitations:
•
http://www.youtube.com/watch?v=ofp-OHIq6Wo
The quantum mechanical model of the atom
• de Broglie: "electrons behave like
particles and have the properties of
waves“
• Heisenberg: "it is impossible to calculate
the exact location of an electron at an
exact moment in time = the uncertainty
principle“
• Schrödinger: "Schrödingers wave
equation describes the region around a
nucleus where there is a high probability
of finding electrons“
http://en.wikipedia.org/wiki/Schr%C3%B6dinger_equation
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According to the QUANTUM MECHANICAL model the
electron has to be treated as both a particle and a
wave.
It is impossible to know exactly where an electron is at
any given time, but it is possibly to calculate the region
in which it is likely to be found.
• Atomic orbital: A region in space around an atomic
nucleus in which there is a high probability of finding
the electron.
Energy levels
• Each electron in atom can be described by four
different quantum numbers:
• The principal quantum number describes the main
energy level, n= 1, 2, 3, etc. Maximum number of
electrons 2n2.
• The second quantum number describes the type of
orbital in that level:
l= 0,1,2,3 or s,p,d,f.
•
The shapes of s, p and d orbitals
/watch?v=sMt5Dcex0kg
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The third quantum number determines the number of
orbitals of each type in each level. For example p x, py,
pz, which are localized along the x, y and z-axes
respectively.
The fourth quantum number describes the spin of the
electron around its own axis, either clockwise or
counter-clockwise, s= +1/2 and -1/2 or ↑↓.
Sub-levels of electrons
Writing electron configurations
• The electron configuration of an element
describes how the electrons are distributed in
sub-levels.
Orbital diagram
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In order to show the spin of each electron, boxes can
be used to represent orbitals.
Aufbau Principle:
• electrons are placed into orbitals of lowest energy
first.
Pauli exclusion principle:
• no two electrons in an atom can have exactly the
same four quantum numbers
no more than two
electrons can occupy any one orbital, and if two
electrons are in the same orbital they must have the
opposite spin = each orbital can contain a maximum
of two electrons
1s
Hund´s rule:
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If more than one orbital in a sub-level is available, for
example px, py, pz , electrons occupy the orbitals single
with parallel spin before they are paired up = orbitals
within the same sub-shell are filled singly first
Ex: Apply the electron-in-a-box to determine the electron configuration
of aluminium.
Al (Z = 13)
Full electron configuration: 1s22s22p63s23p1
Condensed electron configuration: [Ne]3s23p1
Ex: Apply the electron-in-a-box to determine the electron
configuration of zink.
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Electron configurations can be written with the principal
quantum number followed by the letter of the sub-level,
with a subscript to indicate the number of electrons in
that sub-level.
e.g. Na: 1s22s22p63s1
When writing electronic configurations, the sum of the
superscripts must always total the number of electrons
in the atom (or ion).
IB requires that you can write the configuration for
any element or ion up to krypton (Z=36).
Halv-filled and filled sub-levels are more stable:
Cr
Cu
Halv-filled and filled sub-levels are more stable:
Cr
Cu
Electron configuration of ions
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Positive ions are formed by the loss of electrons.
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Negative ions are formed by the gain of electrons.
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(An electron in a doubly occupied orbital is repelled by
its partner and therefore easier to remove than an
electron from a half-filled orbital.)
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Worked example: Apply the electron-in-box method to
determine the electron configuration of the calcium-ion
and the chlorine-ion :
Transition metals
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Elements with partially filled d sub-levels are called
transition metals.
When the d-block elements form positive ions the
energy levels are attracted more strongly to the
nucleus.
The 3d sub-level drops below the 4s sub-level in
energy and therefore the 4s electrons are removed
first.
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Worked example: Apply the in-box method to determine
the electron configuration of the Fe -ion and the
copper-ion :
Evidence from ionization energies
Ionization energy (kJ mol-1)
• The energy required to remove one electron from an atom in
its gaseous state is called the first ionization energy of the
element.
X (g) → X+ (g) + e• The second and third ionization energies would therefore
correspond to the chemical equations:
X+ (g) → X2+ (g) + eX2+ (g) → X3+ (g) + e-
Patterns in first ionization energies
Evidence for sub-levels
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Successive ionization energies for the same element can
also be measured.
Graph of successive ionization energies for sodium.
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As more electrons are removed, the pull of the
positively charged nucleus holds the remaining
electrons more tightly.
Increasingly more energy is required to remove them.
The graph does not increase regularly, which provides
evidence that the main levels are split into sub-levels.
Identify the sub-levels from which the electrons are
removed.
Ionization energies from emission spectra
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The hydrogen emission
spectrum consists of a
series of lines that
converge at higher
energies.
These lines are
associated with electron
transitions from upper
energy levels back down
to lower levels.
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When an electron falls from the limit of convergence
(n=∞) and returns to the ground state (n=1) energy is
given out.
This energy corresponds to the first ionization energy.
Ex. When the hydrogen electron falls from n=∞ to n=1 it
emits energy in the form of light with the wavelength 91,2
nm. This produces a line in the UV region of the
hydrogen emission spectrum (Lyman series).
Calculate the first ionization energy (kJ mol -1) for
hydrogen.
E = hv
c = vλ
E = energy of a photon (J)
h = Planck's constant = 6,63 · 10-34 J s
v = frequency
c = speed of light ≈ 3,00 · 108 m s-1
v = frequency
λ = wavelength (m)