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Atomic Theory
1) An atomic concept of matter is quite old, but there are many modern experiments that support it:
a) Matter can be divided up into pure substances and mixtures
- the former can not be physically separated or broken down into other substances
- mixtures can be made by combining pure substances and can be subsequently separated
again by physical process.
This only makes sense if matter is made of particles.
- pure substances are made of one type of particle, mixtures have two or more types
b) Pure substances can be divided into substances that can be broken down by chemical processes
and those that cannot.
This suggests that some particles are composed elemental particles:
elements
and that other particles are elemental particles that have been joined
together to make a compound particle:
compounds
c) When pure liquids are mixed, the combined volume often is different from the sum of the
original volumes of the separate substances.
This only makes sense if the liquids are made of particles which are of different sizes
(imagine mixing equal volumes of sand and ping-pong balls)
d) When elements are reacted to make compounds (such as copper and nitric acid to make
copper II nitrate) it is possible to subsequently react the compound to get the element back
(as when copper nitrate and aluminum react to make copper and aluminum nitrate).
This only makes sense if matter is made of particles and that the copper particles are not
modified or destroyed, but simply joined to other particles to make the compound.
2) In the 1700s, Antoine Lavoisier carried out chemical experiments in which substances are decomposed
to make other substances. In doing so he:
a) identifies 23 pure substances which cannot be further broken down: elements:
b) makes careful measure of mass in his experiments, thereby discovering that in any chemical
reaction, if care is taken to isolate the reactants and products, that the total mass of the reactants
equals the total mass of the products.
This observation becomes the Law of Conservation of Mass
One inference that can be made from this law is that atoms can neither be created or destroyed.
3) In the early 1800s, Joseph-Louis. Proust observes that:
a) the elements making up compounds exist in a fixed ratio by mass
b) different compounds can be made up of the same elements, but usually each has a different
mass ratio of these elements.
These observations become the laws of Definite Proportions and Multiple Proportions.
The explanation for these laws is that atoms have specific capacities to make connections (bonds) with
other atoms, but that some atoms have different possible capacities for connections. No explanation for
the nature of these connections or why atoms have specific connection capacities is known.
4) At around the same time, Amedeo Avogadro observes that gases combine to form compounds in whole
number ratios by volume. This is eventually seen as indicative of the ratio by number of atoms in which
elements combine to form compounds. In other words, these ratios give us an insight into the chemical
formulas of compounds.
These observations become Avogadro’s Law.
5) Also at around the same time, John Dalton calculates the average atomic masses of several of the
elements by using the mass ratios and the assumed chemical formulas. Since Hydrogen always was
present in the lowest percent by mass in any compound, it was given the mass of 1 u (average atomic mass
unit. The masses of the other elements could be calculated as follows:
Example:
From Avogadro’s Law, water has the formula H2O
The composition of water by mass is 11% Hydrogen and 89% Oxygen
It must be then:
2mHydrogen/mOxygen = 11/89
2(1u)/mOxygen = 11/89
mOxygen = 2(1 u) x 89 / 11
mOxygen = 16 u
In other words, the atomic mass of Oxygen is 16 u.
Dalton took this work and the work of Lavoisier, Proust, and others, and developed a new theory of mater
called the Atomic Theory.
i) Elements are made of extremely small particles called atoms.
ii) Atoms of a given element are identical in size, mass, and other properties;
atoms of different elements differ in size, mass, and other properties.
iii) Atoms cannot be subdivided, created, or destroyed.
iv) Atoms of different elements combine in simple whole-number ratios to form compounds.
v) In chemical reactions, atoms are combined, separated, or rearranged.
Atomic Structure
Starting in the late 1800s, a series of experiments allowed for an understanding of the
internal structure of the atom.
1) Geissler/Crooks/Thompson: The Discovery of Electrons
A glass tube filled with air or another gas will
not conduct electricity. The tube has metal
electrodes at each end which are attached to a
high voltage source of electricity. An ammeter
measures the electric current.
If the air pressure inside the glass tube is
reduced, the gas will conduct electricity and the
gas inside the tube glows. The colour of the
light given off depends on the type of gas. It
does not matter what metal the electrodes are
made out of, the effect is the same.
If the air pressure is reduced almost to zero,
the gas still conducts electricity, but only the
glass itself at the positive end of the tube glows.
If a small paddle wheel is placed in the glass
tube, it rotates when the battery is connected.
In a separate experiment, an object in the tube
cast a shadow in the glow at the positive end of
the tube.
Conclusion
A particle is given off by the negative metal electrode.
It must be a particle because it makes the paddle
wheel rotate. It must be negative because it moves from the negative electrode to the positive electrode
(opposite charges attract and like charges repel). It is easy to remove from the atoms of the electrode
because only electric voltage is needed to remove it. Since the particles are given off by the metal of the
electrodes, they must be a component of the atoms of the metal.
2) Since atoms are neutrally charged, they also must contain a positive particle. Subsequent evidence of this
particle is found in a faint glow at the negative end of gas discharge tubes. These rays were much harder to
deflect with magnetic and electrical fields than the cathode rays (electrons) and thus, were made up of much
more massive particles. Since they were attracted to the negative end of the tube, they must have a positive
charge. The least massive of these were almost 2000 times the mass of the electron and are now known to be
protons.
3) By the early 1900s, people knew that atoms were made of electrons (negative charge) and protons (positive
charge), but it was thought that the electrons and protons were found together as a solid sphere of matter in
an atom. The electrons, being smaller and less massive, were easier to remove.
Thompson was able to show that electrons were a component of all atoms and by studies of how electron rays
were deflected by magnetic and electrical fields that all electrons were the same in terms of charge and
mass. Robert Millikan also found that electrons had a consistent electrical charge in his oil drop experiments.
Michael Faraday was able to show that atoms form ions by gaining or losing specific numbers of electrons.
Arrhenius was able to determine that atoms could form ions in solution by gaining or losing electrons.
4) Radioactivity also had been discovered by this time. Ernst Rutherford conducted exhaustive studies of
radioactivity. He found that there were three types of radiation given off by radioactive material:
alpha particles: heavy particles with a positive charge.
beta particles: light particles with a negative charge.
gamma rays: not a particle, but a type of light with a much higher frequency.
In 1909, Ernest Rutherford devised an experiment to probe the structure of the atom using a beam of alpha
particles given off by a radioactive source. The alpha particles were shot through a thin foil made of Gold.
An alpha particle detector was placed around the gold foil to determine what happened to the alpha particles
after they passed through the Gold foil.
Much to his
amazement, most of
the alpha particles
passed through the
Gold (which is a very
dense material) as if
nothing was there.
However, a few were
deflected off to the
side, and a smaller
number even
bounced back towards the source. This was pretty impressive since alpha particles are very fast moving and
the gold foil was very thin. It was rather like firing a rifle at a sheet of paper and having the bullet bounce
back off the paper.
The conclusions were that atoms must be mostly empty space with a small, very massive core. Since the
positive alpha particles were deflected or repelled by this small core, the core also must have a positive
charge. The stability of atoms suggested that some very strong force holds the protons together in this
nucleus.
The Rutherford model differs from the Thompson model in that:
-
the electrons are separate from the positive core of the atom
electrons orbit around the nucleus
almost all of the mass of the atom is concentrated in a very small nucleus
all of the positive charges are in the nucleus
atoms are mostly empty space with small electrons moving through this space
5) The atomic number was being used as a way to catalogue elements, but is was not known if it was of any
physical significance. In addition, no-one knew how many negative and positive particles the atoms of each
atom had. Also, the reason why the atomic mass of elements increased by the unusual way it did from one
atom to the next was not understood.
Just before the second world war Henry Moseley (a student of Rutherford) set up an experiment to see what
would happen if samples of elements were bombarded with X-rays. He discovered that each element gave off
different X-ray frequencies in response to the bombardment, and that the frequency increased by same
amount each time he went from one atom to the next in the periodic table! The only way that this could
explained is if the number of positive charges in the nucleus increased by one from each element in the
periodic table to the next.
In other words, the atomic number really was a real property of atoms, it was the number of protons in the
nucleus. This conclusion was backed up by several other different experiments, such as the charges that
result when electrons are removed from atoms (if you can only remove so many electrons from a neutral atom,
that number also must equal the number of protons in the nucleus)
6) The Neutron
The number of protons in the nucleus increases by one from one element to the next in the periodic table.
Through mass spectrometry, it was confirmed that the mass increased by more than this. Thus, there must
be another particle in the nucleus, about as massive as a proton, but with no charge: the neutron.
We now have independent evidence of the existence of neutrons. They seem to
help hold all of the positive protons in the nucleus together. Stable nuclei (stable
atoms) have about the same number of neutrons and protons, although larger nuclei
seem to need proportionally more neutrons. Nuclei with either too many neutrons
or not enough neutrons emit radiation as they either fall apart into several smaller
nuclei (nuclear fission) or eject a small number of particles to become more stable.
Given their mass and the fact that some radioactive atoms decay by converting a
neutron to a proton and emit an electron in the process, as is one of the decay
14
mechanisms of Carbon-14:
C —> 14N + e- (Carbon has 6 p+, N has 7)
one might suspect that neutrons are formed by: p+ + e- —> no
7) Isotopes:
It turns out that atoms of the same element do not all have the same mass, thus also must have differing
numbers of neutrons. The atoms of an element with the same number of neutrons are said to belong to the
same isotope. Most elements have a number of stable isotopes, as well as a number of unstable ones that,
none-the-less, decay to other nuclei slowly enough that they are found in significant numbers in nature.
The average atomic mass depends on the relative abundance of each isotope (usually only the stable isotopes
are included in the calculation). These abundances must be calculated from samples collected all over the
earth’s surface.
For example:
element X has two isotopes 40X and 42X
by collecting samples, it is found that 25% of all X are 40X and 75% are 42X
for any given isotope, the mass of the atoms will be close to the mass number.
So, as a rough approximation, the average atomic mass of X will be:
0.25(40 u) + 0.75(42u) = 41.5u
Note: these calculated average atomic masses, as you see them in the periodic table, only apply to the
locations where the samples were taken. If you were a chemist who worked on samples from the earth’s
mantle or on Mars, your periodic table would have slightly different average atomic masses because the
abundances of the different isotopes of each element would be different. Heavier isotopes are just that,
they tend to sink lower into the earth over time than lighter ones. In the formation of the solar system,
lighter atoms tended to get pushed farther from the sun by the solar wind than heavier atoms. Thus, the
proportion of lighter isotopes increases as you get further from the sun. Average atomic masses on Mars
ought to be slightly less for each element than they are on earth.
Even on earth, different environments will have different average atomic masses. For example, since
molecules with heavier masses evaporate more slowly, the oceans tend to get enriched in heavier isotopes over
time.
Atomic nuclei are formed by nuclear fusion which takes place in the cores of stars and in nuclear reactors and
the explosion of nuclear weapons. Very heavy nuclei are formed mostly in super novae, the violent collapse of
the cores of large stars. Any number of protons and neutrons can come together, but only nuclei with the
right ratio of protons and neutrons will tend to stay together.
Atomic Numbers
For any atom: Z is the atomic number
A is the mass number
Z = #p+
N = #no
A=Z+N
The element identity and chemical properties of an atom are determined by Z
The specifications for an atom are written in a standard notation as follows:
“X” is the symbol for the element to which the atom belongs.
for example:
C denotes that the atoms is one of Carbon.
it has 6 protons
N = A - Z, therefore, this atom has 6 neutrons as well.
The charge (C) would be given in the upper-right corner of the symbol, if not zero.
C = #p+ - #eThe mass number for an atom also can be specified as a number following the name of the element:
e.g., Carbon-12 would indicate the same thing as 12C
The average atomic mass is not the same as the mass number for an atom of that element.
i) the proton mass is not the same as the neutron mass
ii) varying amounts of energy are released when protons and neutrons are fused together to make
a nucleus. Since E=mc2, this means that a certain amount of mass is lost in the process.
iii) most importantly, however, the atoms of any element, while they all have the same number of
protons, do not have the same number of neutrons. The given atomic mass actually is an
average atomic mass, the weighted average of the masses of each isotope.
8) The mass spectrometer
The modern method for determining average atomic mass involves using a
device called a mass spectrometer:
It compares the strength of magnetic force needed to accelerate ionized
atoms moving down a glass tube around a bend in the tube to hit a detector
to the force needed to accelerate a reference atom.
Rather than hydrogen, the reference atom used today is 12C
Carbon is less explosive, cheaper, more stable, and easier to contain than
hydrogen. The modern definition for the atomic mass unit is:
1 u = mass of 12C / 12
The average atomic masses in the periodic table are measured in this way.
A sample of atoms of an element are vaporized and fed into the mass
spectrometer. The strength of the magnetic field is adjusted until the
ammeter reads a large current. This means that the ionized atoms are
hitting the detector and drawing an electric current out of the ground. This
magnetic strength is then compared to the one needed to make 12C atoms hit the detector.
The Next Big Question: How are the Electrons Arranged in an Atom?
The evidence to answer this question came from an odd source: light.
Light comes in many colours.
The closer the colour is to the blue end of the spectrum, the more energy it has.
Normal white light (such as sunlight) consists of all colours mixed together. When we split the light into its
component colours with a prism, we see each colour.
The Discovery of Electron Shells
After 1910, it was known that electrons are negative and have little mass. They orbit around a small nucleus
made up of very heavy, positive protons and neutral neutrons. It was not known, however, how the electrons
were arranged around the nucleus. Did they just orbit around the nucleus any old way, or was there some
organization to how they were arranged? In addition, it was known that any charge that was radially
accelerated (such as an electron orbiting a positive charge) would emit light as its velocity changed, thereby
losing energy. Thus, electrons in a Rutherford atom should spiral into the nucleus and the atom collapse into a
Thompson atom. This did not seem to be happening, however.
For 50 or so years previously, it was known that light from a gas discharge tube, when seen through a prism
(or diffraction grating) was found to contain only certain colours from the visible spectrum, unlike sunlight or
the light from incandescent light bulbs which demonstrated the continuous spectrum of the rainbow:
Each element gives off a different set of colours in its spectrum.
Niels Bohr from these observations devised a theory of how electrons are arranged around the nucleus.
Observation 1:
Conclusion:
Protons are positive and electrons are negative, thus they are attracted to each other.
You must add energy to pull electrons away from the nucleus (light, electricity, heat).
When electrons move closer to the nucleus, they must release energy.
the farther they fall towards the nucleus, the more energy they must release.
Observation 2:
Electrons always release energy in the form of light.
The colour of light is directly related to its energy.
Observation 3:
The energy, and therefore colour of light, given off depends on the distance electrons
move towards the nucleus.
Observation 4:
Conclusion:
When you examine the light given off by energized gas, only certain colours of light.
If electrons only give off certain colours, they must only move certain distances.
Main Conclusion
If electrons only move certain distances, they must only exist at certain distances from the nucleus. In other
words, electrons must exist in discrete shells around the nucleus.
This is the first quantum model of the electron: a theory in which the electron can exist in a set of discrete
states, but not in between these states. Quantum models of matter can be described using integer numbers
rather than real numbers. Particles can transition from one state to another, but they do not pass through
any intervening states. If people were like this, we would move from one floor of a building to another, not by
using stairs, but by vanishing from one floor and simultaneously appearing in another. The energy level number
(n) describes what level an electron is in at any time.
The emission spectrum of the Hydrogen atom consists of sets
of lines (mostly outside of the visible spectrum). Each set
is thought to represent the light given off when electrons
transition down into a particular energy level.
i) a) Each shell has a certain capacity to hold electrons.
b) The lowest unoccupied shell that an electron can occupy is called its ground state.
c) An atom is in its ground state when all e- are in the lowest possible shell.
ii) a) Electrons can move up to higher energy levels by absorbing light.
The light frequency must match the change in potential energy of the transition.
b) Electrons are not stable when higher than their ground state and drop back down,
releasing light energy in the process. The frequency of this depends on the amount of
energy released, which depends on how many shells it drops and from which shell.
c) Movement from one shell to another is called a transition.
e- cannot exist between shells (or they could emit any frequency of light).
iii) Experimental evidence can be used to work out the number of e- that each shell
can hold (more on how this can be done soon):
2 in 1st, 8 in 2nd, 8 in 3rd, etc.
Transitions that emit light (in order of increasing frequency):
Transitions that absorb light (in order of increasing frequency):
5, 8, 3
7, 4, 2, 1, 6
Predicting Chemical Behaviour from an Understanding of the Bohr Model
We will work with the simpler of the two types of bonding to start with: ionic.
Sodium and Lithium both will react with Chlorine to make an ionic substance.
In each case, the metal donates its single valence electron to the chlorine to make a positive ion.
The chlorine, accepting this electron, becomes a negative ion and the two ions join together, if the
circumstances are right, to make an ionic compound: sodium or lithium chloride.
We might ask, which of these two metal will react more readily or quickly with the chlorine.
To answer this, consider each step of the reaction.
In the first step, the valence electron of
the metal must be transitioned up from
the ground state to the
outer-most electron shell.
This requires energy to be added to the atom.
Since this step has just about removed the
electron from the atom, ionizing it, we call this
the ionization energy (EI)
Now the electron can shift over to the outermost energy level of the chlorine atom,
This forms the positive (lithium) and negative
(chloride) ions.
Now the electron can drop down to the lowest
available energy level in the chloride ion;
the ground state.
This transition releases energy.
Now we have two stable ions which
can bond together in an ionic bond.
Given what we know about the Bohr diagrams of sodium and lithium, from which atom should it be easiest
(take the least energy) to perform step 1. In other words, which of these two elements should have the
lowest ionization energy? This is the element that should react the fastest or easiest of the two.
The Rutherford Model of the Atom
We will show how Bohr
first shell
second shell
third shell
was able to work out how many electrons each shell could hold:
2
8
8 (actually 18, but we will talk about this next week)
The number of electrons equals the number of protons in a neutral atom.
Electron shells are filled from the first shell outwards.
Thus, each element has a different arrangement of electrons, which can be represented by a Bohr Diagram.
To make a Bohr diagram: i) draw the symbol of the element with the number of protons and neutrons in it.
ii) figure out how many electrons the atom has
iii) draw a circle for the first electron shell
iv) add dots to the circle to represent electrons,
add no more dots that the shell can hold
and add no more dots than the atom actually has electrons
v) subtract the electrons used from the total
vi) if you still have electrons left, add another shell and repeat steps iv to vi
Examples:
If you draw a periodic table of Bohr diagrams, you will see that the elements in the same column have the same electron
state in their outer (valence) shells.
In a neutral atom, the number of protons equals the number of electrons.
How we describe atoms:
Elements: depend on the Atomic number; the atoms of each element have a specific number of protons.
Isotopes: - Not all atoms of an element are exactly alike.
- They all have the same number of protons, but may have different numbers of neutrons
- Atoms of the same element, but different numbers of neutrons are called isotopes.
- Different isotopes have the same chemical properties, but different masses.
- Thus, a heavier atom which behave the same as a lighter atom of the same element, but
substances made out of the heavier atom will be heavier and have slower evaporation rates.
Ions:
if an atom gains or loses electrons, it is no longer neutral, it is charged.
if it gains electrons, it will have a negative charge (more electrons than protons)
if it loses electrons, it will have a positive charge (more protons than electrons)