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Chemistry Vocabulary Review
1. ? – the smallest particle of an element that has all the chemical properties of
that element. Atoms are made of protons, neutrons and electrons.
Ne
H
Ne
Ne
Ne
Neon
Gas
H
H
H
H
H
HH
Ne
Ne
HH
H
H
Hydrogen Gas
HH
2. ? – two or more atoms bonded together.
3. ? – a pure substance made of only one type of atom
H
H
4. ? – a molecule made of more than one type of atom.
It has different kinds of atoms bonded together.
Ammonia
H
NH
H
5. ? – a substance made of more than one type of molecule,
many different molecules not bonded together, and can be separated physically.
For example, a mixture of salt and pepper can be separated by adding water.
The salt will dissolve into the water but not the pepper.
A
strainer or filter can keep the pepper on top while the saltwater goes through.
Heat can then be used to boil off the water leaving the salt in the pan.
It takes too long to use tweezers to separate salt and pepper grains.
A Planetary Model of the Atom / The Bohr model of the Atom
Do you like
my atom?
Before Mr. Bohr’s theory, people didn’t
know how electrons moved around
the nucleus. They thought they all just
orbited in one big cloud like this.
This is the symbol for atomic energy
In 1913 Bohr proposed his quantized shell model of the atom to explain how
electrons can have stable orbits around the nucleus.
The Bohr Model is probably familiar as the "planetary model" of the atom.
In the Bohr Model the neutrons and protons occupy a dense central region called the nucleus,
and the electrons orbit the nucleus much like planets orbiting the Sun (but the orbits of electrons are not
always in the same plane like the planets in the Solar System).
The above image is not to scale. In a real atom the radius of the nucleus is about 100,000 times smaller
than the radius of the entire atom.
Electrons are point particles essentially without a volume.
The electrons take up insignificant amount of space.
Bohr Model for the Hydrogen Atom
The kinetic energy at n=2 would be
greater than n=1. Or else the
electron would spiral closer to the
nucleus. This is because v2 would
provides a nice application of classical orbital mechanics
half as much when n=2
introduces us to some useful concepts inbe
electricity
V = velocity vector of the
electron
to
n=1.
Assuming
you
provides something fun to discuss for thecompared
last few lectures of the term
that we
will not need
to be tested upon.
The fundamental idea of the Bohr model (and models other of the time) was that the electron would orbit the much heavier
spinning around the nucleus.
nucleus in essentially circular orbits. This sounds
very much likeit
the motion
of with
planets aboutyour
the sun, for hand
example, but the
pushed
out
force of gravity is much too weak to account for the energies involved in the atomic spectra and for energies involved in atomic
physics. Instead, the binding force is supplied
by the electromagnetic force, whichto
we study
detail next term. along
This force
perpendicular
itsin more
velocity
operates between charged objects, and, like the gravitational force, is an inverse square law force:
the
radius.
The
force
of
attraction
where Fel is the electrical force between two objects having charges q1 and q2 separated by a distance r. The constant k plays
the role of the universal gravitational constant, and we need not worry too much about it right now. In a hydrogenic atom bechargegreater
is independent
hydrogen, He+, Li++, etc. - those with one would
electron, the nuclear
is +Ze, where Z is(it
the atomic
number (1 for H, 2 for He,
3 for Li, etc.) and e is the 'fundamental charge, 1.6x10^-19 Coulombs. The charge of the electron is -e, so the force on the
of velocity) than the centripetal
electron in a hydrogenic atom is
= Force of
This looks very much like the gravitation formula,
with a slightly
different interpretation
variables involved. In any case,
force
(depends
onof thevelocity).
attain a circular orbit, in the usual way we need equate this central force with the centripetal acceleration. Schematically, the
electrostatic attraction toclassical
orbit of a hydrogenic atom looks like:
The discovery of the nuclear atom, first in the Rutherford laboratory and soon after confirmed by other groups, spawned a
completely new series of theories to explain the discrete lines in the atomic spectra. Initially, the most successful of these was by
Niels Bohr, who was working in the Rutherford laboratory. Bohr devised a model that combined the ideas of Einstein, Planck,
and Rutherford with classical mechanics of orbiting systems that successfully explained the spectra of hydrogenic spectra, i.e.,
the spectra of atoms with one electron. The Bohr model was the first to propose 'quantized states' for atomic systems, a radical
idea at the time. It was soon replaced by 'real' quantum mechanics, but it is still useful to examine the Bohr model here, since it
The Bohr model of a Hydrogen Atom
Centripetal Force -keeps the electron
away from the nucleus
If you push an electron from n=3 to
n=2 then the electron would have
too much velocity to maintain the
orbit at n=2 and the centripetal
force would be greater and it would
Bohr solved these problems by wishing them away in a way that would also predict the spectral properties of hydrogenic atoms.
spin farther from the nucleus back
His postulates were that
into n=3.
the electron could only move in certain non-radiating
states, which he called stationary states. This terminology has been
The force balance condition requires that
There are two apparent flaws with such a model. Firstly, there is no obvious quantization. That is, the final equation does not
quantize r, since we can find the electromagnetic analog a Kepplerian orbits for any r, given the appropriate value of v (and
thus T). A related problem that does not appear (at least not in the same way) in the gravitational case is that the acceleration
of the electron requires that it radiate energy away in the form of light. The mathematics behind this are pretty complex and
you won't see it for a couple of years, but the problem can be easily appreciated. An antenna radiates radio waves by
accelerating charge (i.e., electrons) back and forth along its length. A circular orbit, in this sense, is just an antenna. This
emission of radiation is a dissipative process, as far as mechanical energy is concerned, so classically the electron would spiral
into the nucleus in a very short time.
1 proton in
the nucleus
Z=1
e is the fundamental charge
1.6x10^-19 Coulombs
kept in the 'real' quantum theory.
the electrons could move from one stationary state to another only by making discrete jumps. The radiation emitted is not
related to the electron's motion in either stable orbit, but rather is related by Planck's relationship to the change in energy
in going from one orbit to another:
This means that energy is released
as an electron spins closer to the
That is, one of the Planck/Einstein photons is emitted when an electron moves from a high energy stationary state to a low
energy stationary state.
nucleus. The energy is released in
the angular momentum of the electron in the stationary states was quantized to be an integral number of Planck's
the
form
oforbit
photons
of
light.
constant divided by 2*pi. Since the angular
momentum
of a circular
is mvr, it follows that
the Bohr
quantization
condition requires that
where n is an integer between 1 and infinity. The constant h-bar, defined as h divided by 2*pi, has come to be a more
useful constant than h itself. It is also easier to remember, since its numerical value is very nearly 1x10^-34 J-sec. Thus the
angular momentum of a Bohr orbit is quantized to be an integral number of units of h-bar.
Note that the Planck relation between a photon's energy and the frequency of a light wave
is equivalent to
Thus the use of h-bar rather than h really amounts to finding the angular frequency omega more useful that the cyclic
frequency nu.
How did Bohr arrive at this particular quantization condition? Why does n start at 1 and not zero? Obviously, he guessed.
He certainly tried many others, but this is the one that worked. At the time, there was no real justification.
Force of
electrostatic attraction
= Centripetal Force
keeps the electron away
from the nucleus
Radii of Bohr Orbits
To apply this quantization condition to determine the radii of electronic orbits in atoms, first solve the force balance
equation
for v^2:
and then square the quantization condition
and again solve for v^2:
and set these two equal:
where ao is called the first Bohr radius:
Numerically, ao is equal to 0.0529 nm = 0.529 Angstroms. We see that the angular momentum quantization postulate
Bohr's model explained how the the colors of light were emitted from
hydrogen gas when it was put in a glass tube containing a strong electric
field. (cathode tube) This is known as the emission spectra of hydrogen.
He promptly won the Nobel prize.
The Bohr view of emission and energy is shown schematically in the figure
below. Electrons fall from higher energy orbits to lower ones; resulting in the
emission of photons of energy of different light wave frequencies.
Bohr Atom Pictures for Elements 1 - 13
NAME: _____________
NAME: _____________
Atomic Number: 1
Atomic Mass: 1.0079
Atomic Number: 2
Atomic Mass: 4.003
NAME: _____________
Atomic Number: 3
Atomic Mass: 6.941
protons: _____
neutrons: _____
electrons: _____
protons: _____
neutrons: _____
electrons: _____
protons: _____
neutrons: _____
electrons: _____
NAME: _______________
Atomic Number: 8
Atomic Mass: 16.00
protons: _____
neutrons: _____
electrons: _____
NAME: _______________
Atomic Number: 9
Atomic Mass: 19
protons: _____
neutrons: _____
electrons: _____
NAME: _____________
Atomic Number: 4
Atomic Mass: 9.012
protons: _____
neutrons: _____
electrons: _____
NAME: _____________ per_____
NAME: _____________
NAME: _____________
Atomic Number: 5
Atomic Mass: 10.81
Atomic Number: 6
Atomic Mass: 12.01
protons: _____
neutrons: _____
electrons: _____
protons: _____
neutrons: _____
electrons: _____
Atomic Number: 10
Atomic Mass: 20.18
NAME: _______________
Atomic Number: 11
Atomic Mass: 23
NAME: _______________
Atomic Number: 12
Atomic Mass: 24.31
protons: _____
neutrons: _____
electrons: _____
protons: _____
neutrons: _____
electrons: _____
protons: _____
neutrons: _____
electrons: _____
NAME: _______________
NAME: _____________
Atomic Number: 7
Atomic Mass: 14.01
protons: _____
neutrons: _____
electrons: _____
NAME: _______________
Atomic Number: 13
Atomic Mass: 26.98
protons: _____
neutrons: _____
electrons: _____
Copy this page on back of the worksheet
How to read the Periodic Table:
5
B
Boron
10.81
Atomic Number This is the number of protons = 5
This is also the number of electrons = 5
(only for neutral atoms)
Atomic Symbol
name
Atomic Mass - Atomic Number = 11 - 5 = 6 neutrons
NAME: _B,_ Boron ______
round off 10.81 to 11
Atomic Number: 5
Atomic Mass: 10.81
protons: __5___
neutrons: __6___
electrons: __5___
Draw only up to 8 electrons in the
second
orbital.
3 in
this case =
total
Draw only
up to
2 electrons
in5the
Draw
6 protons
neutronsasashollow
solid circles
first orbital.
Draw
5
circles
Also copy this page on back of the worksheet
The outermost orbital is called the valence
Atoms touch each other at the valence and make
chemical bonds.
Draw the parts of a Hydrogen atom on your worksheet template.
Do these steps on the front of the worksheet
NAME: _ H ,_ Hydrogen
Atomic Number: 1 = number of protons
Atomic Mass: 1.0079 = 1 rounded off
1
1 -1
protons: _____
0 0 calculate the neutrons by subtracting the protons
neutrons: _____
1 the number of electrons is the same as protons
electrons: _____
for a neutral atom.
x
Draw one electron in the first orbital
Draw zero neutrons in the nucleus
Draw one proton in the nucleus
NAME: _ He ,_ Helium
Atomic Number: 2 =
Atomic Mass: 4.003
number of protons
= 4 rounded off
4
2 -2
protons: _____
2 2 calculate the neutrons by subtracting the protons
neutrons: _____
2 the number of electrons is the same as protons
electrons: _____
for a neutral atom.
x
Draw 2 neutrons in the nucleus
Draw 2 protons in the nucleus
x
Draw 2 electrons in the first orbital
NAME: _ Li ,_ Lithium
Atomic Number: 3 =
Atomic Mass: 6.941
number of protons
= 7 rounded off
7
3 -3
protons: _____
4 4 calculate the neutrons by subtracting the protons
neutrons: _____
3 the number of electrons is the same as protons
electrons: _____
for a neutral atom.
x
Draw 1 electrons in the second orbital
x
Draw 4 neutrons in the nucleus
Draw 3 protons in the nucleus
x
Draw 2 electrons in the first orbital
NAME: _ BE, Beryllium
Atomic Number: 4 =
Atomic Mass: 9.012
number of protons
= 9 rounded off
9
4 -4
protons: _____
5 5 calculate the neutrons by subtracting the protons
neutrons: _____
4 the number of electrons is the same as protons
electrons: _____
for a neutral atom.
x
Draw 2 electrons in the second orbital
x
Draw 5 neutrons in the nucleus
Draw 4 protons in the nucleus
x
x
Draw 2 electrons in the first orbital
NAME: _ B, Boron
Atomic Number: 5 =
Atomic Mass: 10.81
number of protons
= 11 rounded off
11
5 -5
protons: _____
6 6 calculate the neutrons by subtracting the protons
neutrons: _____
5 the number of electrons is the same as protons
electrons: _____
for a neutral atom.
x
Draw 3 electrons in the second orbital
x
Draw 6 neutrons in the nucleus
Draw 5 protons in the nucleus
x
x
x
Draw 2 electrons in the first orbital
NAME: _ C, Carbon
Atomic Number: 6 =
Atomic Mass: 12.01
number of protons
= 12 rounded off
12
6 -6
protons: _____
6 6 calculate the neutrons by subtracting the protons
neutrons: _____
6 the number of electrons is the same as protons
electrons: _____
for a neutral atom.
Draw 4 electrons in the second orbital
x
x
x
x
x
x
Draw 6 neutrons in the nucleus
Draw 6 protons in the nucleus
Draw 2 electrons in the first orbital
NAME: _ N, Nitrogen
Atomic Number: 7 =
Atomic Mass: 14.01
number of protons
= 14 rounded off
14
7 -7
protons: _____
7 7 calculate the neutrons by subtracting the protons
neutrons: _____
7 the number of electrons is the same as protons
electrons: _____
for a neutral atom.
Draw 5 electrons in the second orbital
x x
x
x
x
x
x
Draw 7 neutrons in the nucleus
Draw 7 protons in the nucleus
Draw 2 electrons in the first orbital
NAME: _ O, Oxygen
Atomic Number: 8 =
Atomic Mass: 16.00
number of protons
= 16 rounded off
16
8 -8
protons: _____
8 8 calculate the neutrons by subtracting the protons
neutrons: _____
8 the number of electrons is the same as protons
electrons: _____
for a neutral atom.
Draw 6 electrons in the second orbital
l l
l
l
l
l
l l
Draw 8 neutrons in the nucleus
Draw 8 protons in the nucleus
Draw 2 electrons in the first orbital
NAME: _ F, Fluorine
= number of protons
= 19 rounded off
19
9 -9
protons: _____
10 10 calculate the neutrons by subtracting the protons
neutrons: _____
9 the number of electrons is the same as protons
electrons: _____
for a neutral atom.
Atomic Number: 9
Atomic Mass: 19
Draw 7 electrons in the second orbital
l l
l
l
l
l
l
l l
Draw 10 neutrons in the nucleus
Draw 9 protons in the nucleus
Draw 2 electrons in the first orbital
NAME: _ Ne, Neon
Atomic Number: 10 =
Atomic Mass: 20.18
number of protons
= 20 rounded off
20
10 -10
protons: _____
10 10 calculate the neutrons by subtracting the protons
neutrons: _____
10 the number of electrons is the same as protons
electrons: _____
for a neutral atom.
Draw 8 electrons in the second orbital
l l
l
l
l
l
l
l
l l
Draw 10 neutrons in the nucleus
Draw 10 protons in the nucleus
Draw 2 electrons in the first orbital
NAME: _ Na, Sodium
Atomic Number: 11 =
Atomic Mass: 23
number of protons
= 23 rounded off
23
11 -11
protons: _____
12 12 calculate the neutrons by subtracting the protons
neutrons: _____
11 the number of electrons is the same as protons
electrons: _____
for a neutral atom.
Draw 8 electrons in the second orbital
l l
l
l
l
l
l
l
l
ll
Draw 12 neutrons in the nucleus
Draw 11 protons in the nucleus
Draw 2 electrons in the first orbital
Draw 1 electron in the third orbital
NAME: _ Mg, Magnesium
Atomic Number: 12 =
Atomic Mass: 24.31
number of protons
= 24 rounded off
24
12 -12
protons: _____
12 12 calculate the neutrons by subtracting the protons
neutrons: _____
12 the number of electrons is the same as protons
electrons: _____
for a neutral atom.
l
Draw 8 electrons in the second orbital
l l
l
l
l
l
l
l
l
ll
Draw 13 neutrons in the nucleus
Draw 12 protons in the nucleus
Draw 2 electrons in the first orbital
Draw 2 electrons in the third orbital
NAME: _ Al, Aluminum
Atomic Number: 13 =
Atomic Mass: 26.98
number of protons
= 27 rounded off
27
13 -13
protons: _____
14 14 calculate the neutrons by subtracting the protons
neutrons: _____
13 the number of electrons is the same as protons
electrons: _____
for a neutral atom.
l
Draw 8 electrons in the second orbital
l l
l
l
l
l
l
Draw 2 electrons in the first orbital
l
l
ll
Draw 14 neutrons in the nucleus
Draw 13 protons in the nucleus
l
Draw 3 electrons in the third orbital