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Atomic Concepts
Democritus (460 – 370 BC) – indivisible tiny
particles called atoms
Plato and Aristotle – atoms cannot be indivisible
particles
Issac Newton (1642 – 1727) – atoms
John Dalton (1803 – 1807) – Dalton’s atomic
theory
 elements are composed of atoms
 all atoms of a given element are identical to
one another but different from atoms of all other
elements
 atoms of one element cannot be changed
into another, created or destroyed
 compounds are formed when atoms combine
in new ways
Dalton’s atomic theory explained
- law of constant composition – in compound,
number and kinds of atoms are constant
- law of conservation of mass - mass present
after a chemical reaction is the same as total mass
present before the reaction
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James Maxwell (1870) – mathematical description
of the general behavior of light, described how
energy in the form of radiation can travel through
space as electric and magnetic fields
J.J. Thomson (1897) – plum pudding model –
negative charged particles swimming in a sea of
positive charge, build a cathode ray tube with a
metal cylinder on the end, cylinder had two slits in it,
leading to electrometers, which could measure small
electric charges, by applying a magnetic field across
the tube, there was no activity recorded by the
electrometers and so the charge had been bent away
by the magnet, proved that the negative charge and
the ray were inseparable and intertwined, also
proved that cathode rays carried a negative charge
because rays were deflected by an electric field
Robert Millikan (1909) – measured charge of an
electron, oil-drop experiment, oil which had picked
up extra electrons were allowed to fall between two
electrically charged plates, calculated the charge on
the drops by measuring how the voltage on the
plates affected their fall, charges were always
integral multiples of 1.602 x 10-19 C
Ernest Rutherford (1911) – 3 types of radiation
(alpha, beta and gamma), gold foil experiment - shot
alpha particles at gold foil and found most of the
alpha particles passed through the foil, discovered
protons, stated atom has a small dense positive
nucleus and that most of the atom was empty space
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Bohr atoms (1913) – small dense positively
charged nucleus surrounded by electrons in specific
energy levels
Louis de Broglie (1924) – all matter has both wave
and particle characteristics
James Chadwick (1932) – discovered neutons
Albert Einstein (1940) – photoelectric effect –
when light with certain frequencies strikes a piece
of metal, it emits electrons from the metal, radiant
energy behaves as a stream of tiny packets of
energy called photons (have properties of waves)
Aufbau Principle – Electrons are added in order of
increasing energy: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14
Heisnberg Uncertainty – it is not possible to know
both the position and momentum of an electron at a
particular moment, electron orbitals are described in
terms of probability
Hund’s Rule – electrons will enter empty orbitals
of equal energy when they are available
Pauli Exclusion Principle – no two electrons in an
atom have the same set of four quantum numbers
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Max Planck – father of quantum physics, energy
can only be emitted or absorbed from atoms in fixed
amounts (quantum)
quantum - the minimum amount of energy that
can be gained or lost by an atom
energy emitted by hot objects is quantized
Planck’s constant (h) – 6.626 x 10-34 J . s
Erin Schrӧdinger – apply probability to describing
the volume of space where an electron would be
located
charge electron
charge proton
-1.602 x 10-19 C
+1.602 x 10-19 C
protons = electrons in an atom
1 amu = 1.66054 x 1024 g
1 g = 6.02214 x 1023 amu
proton 1.0073 amu
neutron 1.0087 amu
electron 5.486 x 10-4 amu
A–Z=N
atomic mass – atomic number = number of
neutrons
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stability of nucleus – use neutron/proton ratio, 1:1
stable
elements atomic number lower than 20 are light
elements and have a nuclei ratio of 1:1 (same
amount of protons and neutrons)
elements atomic numbers 20 to 83 are heavy
elements and nuclei ratio is not 1:1, the reason is
because of the repulsive force between protons, the
stronger the repulsive force, the more neutrons are
needed to stabilize the nuclei
Belt of stability
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mass spectrophotometer – atoms can be
deflected by magnetic fields provided the atom is
first turned into an ion, electrically charged particles
affected by magnetic field and neutral ones aren’t
stages in mass spectrophotometer
 ionization – atom ionized by knocking one or
more electrons off to give a positive ion
 acceleration – ions accelerated so they all
have the same kinetic energy
 deflection – ions deflected by magnetic field
according to their masses, lighter ones more
deflected and the more the ion is charged, the
more it gets deflected
 detection – beam of ions detected
electrically
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isotopes – same number of protons but different
number of neutrons
11
C
13
C
14
C
calculate atomic mass of an element from its
isotopic abundances
3 isotopes of silicon
92.23% of 28Si with atomic mass of
27.97693 amu
4.68% of 29Si with atomic mass of
28.97649 amu
3.09% of 30Si with atomic mass of
29.97377 amu
27.97693 amu (0.9223) + 28.97649 amu (0.0468)
+ 29.97377 amu (0.0309) = 28.09 amu
electromagnetic radiation – waves produced by
the motion of electrically charged particles, travel
through empty space as well as air and other
substances, can act like waves, stream of particles,
called photons, that have no mass, photons with the
highest energy correspond to the shortest
wavelength
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matter can gain or lose energy in small, specific
amounts called quanta
quantum - the minimum amount of energy that
can be gained or lost by an atom
energy emitted by hot objects is quantized
photoelectric effect - when electrons are emitted
from metal’s surface when light of a certain
frequency shines on
de Broglie equation – λ = h / m . v
λ - wavelength
momentum – p = m . v
energy of an electron – En = -2.178 x 10-18 / n2 J
J – joules
n – principal quantum number
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speed of light (c) – 2.998 x 108 m / s
(same as
ms-1)
diamagnetic – all subshells are filled with
electrons
paramagnetic – subshells are not completely filled
with electrons
Electronic Structure of Atoms
n
(principal
E level)
1
l (azimuthal,
shape)
ml (magnetic)
ms (spin)
0
0
+/- 1/2
2
0
1
0
-1, 0 , +1
+/- 1/2
+/- 1/2
3
0
1
2
0
-1, 0 , +1
-2, -1, 0, +1, +2
+/- 1/2
+/- 1/2
+/- 1/2
4
0
1
2
3
0
-1, 0 , +1
2, -1, 0, +1, +2
-3, -2, -1, 0, +1,
+2, +3
+/+/+/+/-
1/2
1/2
1/2
1/2