Download 2 Atomic Theory Development of Theory Historical Atomic Models

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Atomic Theory
Development of Theory
Historical Atomic Models
Democritus
Greek philosopher who postulated that matter is
comprised of atoms as the smallest part (ca 400 BC)
atoms are the smallest, indivisible part of an element,
are alike for one kind of element, combine
chemically in whole ratios, and compounds are
composed of elements (1803)
John Dalton
Max Planck
first postulated that energy is quantized, originator of
quantum theory (1900)
J.J. Thompson
“plum-pudding” model – negative electrons (plums)
are located in a positively charged pudding (1903)
Hantaro Nagaoka
“Saturnian” model – large nucleus with electrons
orbiting in rings (1904)
Robert Millikan
measured the charge and mass of an electron (1908)
Ernest Rutherford
atom has a small, positive, central nucleus, which
contains the mass and is surrounded by a cloud of
negative electrons [correct model] (1911)
Discovered the proton. (1917)
“planetary” model – the nucleus is surrounded by
electrons orbiting in rings (1913)
Neils Bohr
Edwin Schroedinger mathematical wave equation led to prediction of the
possible states for an electron [correct, part of
orbital theory] (1926)
Rutherford Experiment
Werner Heisenberg
“uncertainty” principle correctly states that it is
impossible to predict the exact position and
momentum of an electron (1927)
James Chadwick
discovered the neutron which served as a catalyst for
the growth of nuclear physics (1932)
method:
involved shooting alpha particles (He2+) at a sheet of gold foil
postulate:
the deflection of the alpha particles would determine the
location of the mass within the atom
most particles went straight through, while some deflected
back
results:
conclusions: atom is mostly empty space {why many particles went
straight}, with almost all the mass in a small positively
charged nucleus {why some particles were deflected
backwards}
Nucleus
Central portion of the atom, positively charged and contains the mass .
Electron cloud
The bulk of the space of an atom, mostly empty, negatively charged.
Basic Subatomic Particles
electron
proton
negative charge (–)
positive charge (+)
located in electron cloud
located in nucleus
0u
1u
neutron
neutral
located in nucleus
1u
( )
Note that for an individual atom, the number of protons and neutrons never changes in ordinary reactions.
The number of electrons can change, which effects the charge of the atom, but the nucleus does not change.
Atomic Mass, Y
the sum of the protons and neutrons. p + n
e.g. an atom with 6 protons and 7 neutrons has an atomic mass of 13
Atomic Number, Z
The number of protons. This defines the element. For example carbon
always has 6 protons, but is known to have 6 neutrons (Y=12) or 7 neutrons
(Y=13) or 8 neutrons (Y=14) See isotopes below.
Isotope
Atoms with the same number of protons (same element), but with different
number of neutrons.
Percent Abundance
the percentage of one isotope for an element
Average Atomic Mass, Yavg
a weighted average of all known isotopic masses for an element
Yavg
Y1
Y2

where X = percent abundance as a decimal
Y1 and Y2 are isotopic masses
Charge
Some particles emit an electric force creating a field of force around the
particle. This field attracts opposite fields while repelling similar fields.
These fields are positive (+) or negative (-). The absence of the field is
neutral ( ).
Ion
a charged atom or molecule
Cation
positive ion, lost electrons
Anion
negative ion, gained electrons
To calculate charge
atom – number of excess protons or electrons. e.g. if an atom contains 6
protons and 7 electrons, thus 6 + charges and 7- charges with a net of 1charge
molecule – the sum of the oxidation numbers for each atom is the charge,
e.g. sulfate is comprised of one S at a +6 oxidation number and four O’s
each at a -2 oxidation number, thus sulfate has a 2- charge because
1S + 4O’s = (+6) + 4(-2) = -2
Oxidation Number
the apparent charge of an atom in the molecule. Some oxidation numbers
can often be found from the atom’s location on the periodic table, Group 1
is +1, Group 2 is +2, H is +1 (or -1 in hydrides), O is -2 (-1 is peroxides), in
binary ionic compounds the halogens are -1. Otherwise the oxidation
number is calculated .
example: given, NaClO4, where Na = +1 (Group 1), O = -2,
since Na + Cl + 4 O = 0,
then Cl = +7.
Note that for a single atom the charge is the oxidation number.
Oxidation
the loss of electrons in a chemical change; increase in oxidation number
Reduction
the gain of electrons in a chemical change; decrease in oxidation number
Note that oxidation cannot happen without reduction and vice versa.
Electrons
Electron Spin
from probability, electrons are said to spin up (↿) or spin down (⇂).
Electron Pair
energy.
(↿⇂) - combination of a spin up (↿) with a spin down (⇂). Pairing requires
Valence electrons
electrons in outermost energy level. These are the electrons involved in
bonding and reactions.
Aufbau Principle
lowest energy orbitals fill first with electrons (fill diagram from the bottom
– up, lower energy state is preferred)
Hund’s Rule of Multiplicity
if two or more orbitals of equal energy are available, electrons will occupy
them singly before filling them by pairing. Electron pairing requires
energy, the lower energy state is always preferred so the electrons stay
single. Following the axiom that the lower energy state is preferred the
electrons will pair when the choice is pairing or moving to an orbital that is
higher in energy
Paule Exclusion Principle
no two identical electrons can occupy the same orbital – means that only
electrons of opposite spin may be in the same orbital
Energy Level, n
A discrete distance from the nucleus. The further the distance, the greater
the energy needed for an electron to be located there.
Electrons must gain a set amount of energy to move to a higher energy level
further from the nucleus (the set amount of energy is said to be quantized).
The energy is released as a set amount of energy if the electron moves
closer to the nucleus to a lower level.
Ground State
All electrons in the lowest energy level possible.
Excited State
One or more electrons not at ground state.
Orbitals
A defined region (shape) of space, where it is most probable to find an
electron. Each orbital contains 0, 1, or 2 e’s. There are four classes of
orbitals: s, p d, f. Each class of orbital can have certain types. For instance
the p-orbital has 3 types: px, py, pz. Each orbital is hourglass shaped and is
aligned along an axis of space.
s: 1 type, total of 2 e’s, 1 pr
p: 3 types, total of 6 e’s, 3 pr’s
d:
Sublevel, l
Orbital type, m
5 types, total of 10 e’s, 5 pr’s
f:
7 types, total of 14 e’s, 7 prs
Indicates the class of orbital (s, p, d or f) present in the energy level.
l = 0, …(n-1)
l
Electron spin, ms
Indicates the specific orbital (e.g. px, py, pz). m l = - l, …, + l
Electrons spin up (↑), ms = 12 , or down (↓ ) ms = - 12
(n)s
(n)p
(n-1)d
(n-2)f
Electron configuration
states the arrangement of electrons within the electron cloud; includes the
energy level, orbital type and number of electrons.
examples: H = 1s1
N = 1s2 2s2 2p3
Notes -
3
All families have the same valence electron configuration
noble gas configuration
ns2np6
halogen configuration
ns2np5
chalcogen (O-family) configuration
ns2np4
Atom Stability
Nuclear
Radioactivity
the release of energy and/or particles resulting from an unstable nucleus
There is no set rule for stability, but from experiment stability is based on
the neutron to proton ration, the further the value of np is from 1, the more
likely the isotope is radioactive.
12
6
13
6
C
n
p
C
= 66
7
6
14
6
C
8
6
furthest from 1, so most likely radioacitve
Nuclear Transformations
a change in the number of protons and /or neutrons in the nucleus as a result
of radioactive decay
Half-life
The time it takes for half of a sample of a radioactive isotope to decay. For
example, the half-life of 32P is 14 days. So after 14 days a 50 g sample of
32
P is now 25 g of 32P and 25 g of 32S. (see beta decay below)
Types of radioactive decay
4
o , Alpha Particle, 2 He positive He nucleus ejected from the nucleus ,
o
Beta Decay,
0
1
o , Gamma Rays
32
15
32
P  10 e + 16
S
high energy photon emitted as nucleus moves from excited to lower energy
state
o EC-electron capture
Fr  24 He + 219
85 At
high energy e- is ejected from the nucleus (n p + e-),
e
o Positron Emission,
223
87
0
1
232
90
Th *  132
90Th +
(*=excited state)
e positive particle ejected from nucleus (p 01 n +
0
1
e ),
30
15
30
P  10 e + 14
Si
e- falls into nucleus combining with a proton and forming a neutron,
202
81
Tl + 10 e  202
80 Hg
4
Periodic Table
History
Dmitri Mendeleev
Henry Mosely
Periodic Law
Wrote the 1st periodic table based on increasing atomic mass and similar
properties.
Left gaps where necessary in order to line-up families with similar
properties.
Predicted products of missing elements that, when discovered, would fill-in
the gaps
Created the modern periodic table based on increasing atomic number
The physical and chemical properties of the elements are periodic functions
of their atomic number.
Layout
Period
Group/Family
Trends
Electron shielding
Horizontal rows
A period is likened to an energy level when completing energy level
diagrams.
Moving left to right, the effective nuclear charge (the attraction between the
valence electrons and the nucleus) increases, this causes the atomic radius to
decrease, and electronegativity and ionization energy to increase.
A vertical column
Elements in the same family have the same valence e-config, and thus
similar properties
When moving down a group the distance (# of energy levels) between the
nucleus and the valence electrons increases causing the attraction between
them to decrease, so atomic radius increases down a group while the
electronegativity and ionization energy decrease.
the masking of the nucleus by the kernel electrons. Shielding is constant
within a period, but grows down a group
Effective nuclear charge
the charge felt by each valence electron.
Calculated by protons – kernel electrons
Increases left to right across a period, but is constant in a group
Electronegativity
the ability to attract electrons in a covalent bond
trend =
↑
First Ionization Energy
the energy needed to remove one electron
trend =
↑
Atomic Radius
distance from the nucleus to the valence energy level
trend = ↓
examples:
Which is more electronegative, K or Cl?
ans = Cl
Which has the larger atomic radius, S or As?
ans = As