Download Unit 3 – Atomic Theory

UNIT 3 – ATOMIC THEORY
CH1030
Mark Stacey
ATOMIC THEORY
A molecule is the smallest piece of a substance that retains the
properties of that substance. For example, a water molecule (H2O)
acts like water but if it is broken down further (into lone H or O) it no
longer behaves as water.
An atom is the smallest piece of a single element that still retains the
properties of that element. A molecule can contain one or many atoms.
Atoms are effectively the smallest whole units of matter. While atoms
themselves are made of smaller parts, these parts do not exist
independently in most everyday life situations.
ATOMIC THEORY
Atoms were first theorized by the ancient Greeks as a term to
describe the smallest piece of a substance that still acted as that
substance. However, this was long, long before an individual atom was
very observed.
This is because atoms are extremely tiny, in the range of
0.0000000001m in diameter. This means that even a few milligrams
of a substance contains many, many atoms.
The periodic table shows the ~120 types of atoms we have observed
or predicted to exist. All matter is made up of one or a combination
of these atoms.
ATOMIC THEORY
We describe different substances using chemical formulas, a
shorthand method to describe the types and quantities of atoms in a
single molecule.
All types of atoms have one or two letter shorthand. The first letter is
always capitalized and the second (if present) always is lowercase.
Some elements are named based on older or non-English names and
so their shorthand name may not align well with their modern English
full name.




Oxygen – O
Carbon – C
Sodium – Na (from Latin: natrium)
Iron – Fe (from Latin: ferrum)
ATOMIC THEORY
An element is a chemical containing only one type of atom. Elements
may exist as single atoms, or a grouped molecules – the important
thing is that all the atoms are the same type.
Compounds contain two or more types of atoms within their individual
molecules. Simply having two types of atoms floating around
separately is not a compound – that is a mixture.
Mixtures contain two or more types of molecules. These molecules
themselves can be elements or compounds.
ATOMIC THEORY
ATOMIC THEORY
(Pure)
Element
(Pure)
Compound
Mixture
Types of Atoms
1
2+
2+
Types of
Molecules
1
1
2+
Number of Atoms
per Molecule
1+
2+
n/a
SUB-ATOMIC PARTICLES
Atoms may be the smallest functional units of matter, but they are not
the smallest things to exist. Atoms are made of three major subatomic particles:
Neutrons (n) make up the core of an atom called a nucleus. These
particles have no charge to them and are largely only contribute to
the size and mass of an atom.
Protons (p+) are also found within the nucleus. These have a similar
size and mass to neutrons, but also have a positive charge. This makes
the nucleus as a whole have a net positive charge.
SUB-ATOMIC PARTICLES
Electrons (e-) do not exist within
the nucleus of the atom. Instead,
they orbit the nucleus from a
relatively far distance. Electrons
are much, much smaller and
lighter than protons or neutrons.
Electrons have negative charge,
equal and opposite to a proton’s
positive charge.
Electrons are kept in their orbits
partially by this positive-negative
attraction.
A representation of an atom using the
“planetary model”. Note that the sizes
and distances shown here are not
realistic. Electrons are significantly
smaller than protons or neutrons, and
the distance from electrons to the
nucleus is significantly larger.
SUB-ATOMIC PARTICLES
As electrons are more “accessible” to the outside, it is electrons that
are more readily lost, gained or shared to other atoms. Most chemical
changes involve the redistribution of electrons between atoms.
The nucleus is significantly more stable. However, nuclei can be joined
together or split apart. These types of reactions take tremendously
more effort to accomplish compared to regular chemical reactions
involving electrons. As these reactions involve the nucleus, these are
called nuclear reactions.
SUB-ATOMIC PARTICLES
Sub-Atomic
Particle
Neutron
Location in
Atom
Nucleus
Charge
None
Size
10-15 m
Mass
Reactivity
10-24 g
Nuclear
Reactions
Only
Proton
Nucleus
Positive
10-15 m
10-24 g
Nuclear
Reactions
Only
Electron
Orbitals
Around
Nucleus
Negative
10-16m
10-28 g
Any Chemical
Reaction
MODELS OF THE ATOM
While the concept of the atom has existed for two thousand years, the
actual makeup of atoms themselves has only been explored in the
past hundred years.
The Greek idea of the atom as a single, whole, indivisible particle was
codified into what is called the Dalton model of the atom.
Under the Dalton model:




The smallest pieces of an element possible are atoms
Every element has a different atom unique to it
Atoms cannot be further divided
Atoms are the smallest particles to exist
MODELS OF THE ATOM
At the end of the 19th century, the electron was discovered to exist. As
well, the electron was observed to be something that could be isolated
from all elements, meaning that all atoms had to be partially made up
of these electrons. Electrons were discovered to be negative in charge.
As atoms are neutral in charge, scientists deduced that there must also
be a positively-charged species in the atom to balance out the
electron. The proton was deduced even before it was isolated.
This created the Thompson model:
 Atoms are no longer singular pieces, but instead contain protons and electrons (and
potentially other yet-to-be-discovered parts
 Electrons and protons are non-specifically arranged in the “mixture” that is an atom
 Thompson likened the atom to a plum pudding, were electron and protons are
suspended in the atom, like the plums in the pudding.
MODELS OF THE ATOM
MODELS OF THE ATOM
Approximately 20 years later, experiments were done to prove the
existence of the nucleus. This was done by sending a beam of
radiation through an atom that would only react to protons and
observing the “shadow” cast from the protons.
If Thompson’s model were true, the “shadow” would appear as a
somewhat random distribution, as the protons should have no
organization.
Instead, it was observed that the majority of the atom cast no
“shadow” at all, and only a very small area at the centre did. This
lead to the Rutherford model:
 Protons are densely-packed into a central nucleus
 Protons only exist within this nucleus.
MODELS OF THE ATOM
MODELS OF THE ATOM
Not long after, the neutron was discovered to exist. It’s lack of a
charge made it harder to observe than the charged particles.
The neutron was discovered to also exist with in the nucleus along with
the proton.
As time went on, the Bohr model would be developed, which
explained that electrons not only orbited the nucleus randomly, but in
a very specific series of orbits. This is sometimes known as the
planetary model, as it resembles a star and its orbiting planets.
Recently, the idea of electrons forming discrete circular orbits has
been replaced with the idea that electrons exist within an orbital
cloud. This is due to the fact that objects as small as an electron can
behave differently than larger objects. This is the quantum model.
MODELS OF THE ATOM
ATOMS AND ELEMENTS
All atoms are a combination of protons, neutrons and electrons. What
makes one atom different from another (and therefore what makes
one element different from another) is the numbers of protons,
neutrons and electrons.
As protons make up the nucleus and have a charge, we define each
element by the number of protons. Neutrons also make up the
nucleus, but as they are neutrally-charged, they do not play into the
behaviors of the atom significantly.
The number of protons in an atom is known as the atomic number (Z).
The simplest element, hydrogen has one proton (Z=1), while the
heaviest elements currently known have over 110 protons (Z=110+)
ATOMS AND ELEMENTS
The atomic number is the large
number displayed with an
element on the periodic table.
Atomic numbers are always
whole numbers (no decimals), as
there cannot be a fraction of a
proton.
Atomic numbers cannot be less
than 1. All atoms have at least
one proton.
ATOMS AND ELEMENTS
The number of protons defines which element an atom is. For example,
an atom with 8 protons is oxygen, regardless of the number of
neutrons and electrons. If there are 8 protons, then it is some “variant”
on oxygen.
As neutrons have no charge, they effectively only contribute mass and
size to the nucleus. Therefore if a oxygen atom with 10 neutrons is not
terribly different from an oxygen with 8 neutrons, other than being
heavier.
ATOMS AND ELEMENTS
The various versions of an atom containing differing numbers of
neutrons are called isotopes. Isotopes generally have the same
chemical properties (as reactivity is more influenced by
protons/electrons).
Generally, most atoms have a very similar number of neutrons and
protons. For example, the common isotopes of carbon (Z=6) contain 6,
7, or 8 neutrons. The “default” configuration is an equal number of
neutrons and protons, but this does not have to be the most common
version in nature.
The mass number (A) of an atom is the sum of its protons and
neutrons. Electrons are not counted in this as their weight is so much
less than protons and neutrons that it is insignificant.
ATOMS AND ELEMENTS
ATOMS AND ELEMENTS
The number of neutrons can be found by subtracting the atomic
number from the mass number:
 Neutrons = A – Z
Isotopes are indicated by writing the mass number as a superscript
before the chemical symbol, or by writing it after symbol with a
hyphen.
Carbon with A=13:
 13C or C-13
 This is pronounced “carbon thirteen”
ATOMS AND ELEMENTS
Hydrogen is the smallest element,
with only one proton (Z=1).
Hydrogen is also the only
element that can exist without
any neutrons.
1H
has no neutrons:
N=A–Z
N=1–1
N=0
Hydrogen-1, commonly
known as protium.
ATOMS AND ELEMENTS
Other isotopes of hydrogen
have neutrons, however:
H-2 (common name: deuterium)
has 1 neutron:
 N=A–Z
 N=2–1
 N=1
H-3 (common name: tritium) has
2 neutrons:
 N=A–Z
 N=3–1
 N=2
Water molecules made up of H-2 rather
than H-1 are called “heavy water” as the
extra neutron per hydrogen atom increases
the mass of the water. Heavy water has
many industrial uses.
RADIOACTIVITY
Not all isotopes of an element
are stable. Some isotopes are
prone to decay into smaller
pieces (either casting off some
sub-atomic particles and/or
splitting entirely into two or
more nuclei).
This is a nuclear reaction known
as radioactive decay.
A depiction of radioactive decay. In this
image, an alpha particle (2 neutrons
and 2 protons) is ejected from a larger
nucleus.
RADIOACTIVITY
One useful type of radioactive
decay is nuclear fission
(fission = splitting). In this
reaction, certain specific
elements have their nuclei
broken down into smaller parts.
This reaction releases a
tremendous amount of energy,
which can be used for an
explosion (nuclear weaponry),
or to power and electric
generator (nuclear reactor).
U-235 is bombarded with a neutron to
create U-236, which rapidly undergoes
nuclear fission to yield Kr-92, Ba-141
and 3 lone neutrons. In many fission
reactions, these by-products will
themselves undergo fission until they
reach a stable atomic isotope.
RADIOACTIVITY
Several isotopes undergo radioactive decay naturally at extremely
steady rates and over very long periods of time.
By knowing the original amount of a decaying isotope that was
present and comparing it to the current amount, radiometric dating
can be done to estimate the age of an item.
Carbon-14 takes 60,000 years to fully decay. Carbon-14 dating
allows scientists to date any item from within this time period.
Naturally, items older than this cannot be C-14 dated.
For older items, such as dinosaur fossils, other elements such as
Uranium-238 are used, as their decay occurs over a much longer
timeframe. U-238 dating is highly accurate, allowing us to measure
things up to 4.5 billion years old to within 1% of error.
RADIOACTIVITY
Radiation exposure is very harmful to living things.
The energy and small particles released from radioactive reactions
will disrupt and destroy the biological molecules that make up living
cells. This in turn will cause severe disruption of life processes,
localized cell death, or complete death depending on the exposure.
Radiation shielding uses materials (such as thick lead plates) that
absorb the energy and particles released from radioactive processes.
RADIOACTIVITY
However, many radioactive isotopes have beneficial uses in medicine:
Radioactive (but extremely harmless) technetium-99 can be injected
into the body. As it circulates, radiation detectors can be used to
observe its location and flow – helping to map out the body and
detect circulation problems.
Radioactively-tagged molecules (a otherwise-normal molecule
containing a radioactive atom) are often used in research to find out
what happens to molecules in living things. For example, a protein may
be made containing radioactive sulfur and fed to an organism. Later,
the scientists examine the organism and see what chemicals that
organism has incorporated that sulfur into, thus tracking what
substances that organism transforms its dietary protein into.
RADIOACTIVITY
Radiation therapy uses the destructive powers of radiation in as
careful and specific a manner as possible.
Radiation is used on extremely-specific regions of the body to kill
cancer cells. Naturally, the radiation cannot differentiate between
cancer and healthy cells and thus radiation therapy must be applied
to as precise an area as possible to reduce the amount of healthy
cells damaged.
FUSION
The opposite reaction to
radioactive decay is that of
nuclear fusion. Fusion, as the
name implies, involves the
combination of nuclei into a new,
larger nucleus.
Both nuclei are positivelycharged and thus will resist
being pushed together. This
makes the energy needed for
fusion extremely high.
Thermonuclear weapons create
a small fusion reaction to unleash
tremendous destructive power.
Nuclear fusion of H-2 and H-3 into He-4
and a lone neutron. This reaction yields
a large amount of energy (measured
here in electronvolts (MeV)).
FUSION
After the big bang, the
majority element was (and still
is) hydrogen.
Larger elements have been
formed over time within stars.
The immense size and pressure
of a star’s core causes fusion to
occur, creating larger and
larger elements, depending on
the size of the star.
Larger elements are only found
in regions of our universe that
have had stars in them at some
point.
A depiction of the inner portion of a star
showing subsequently heavier elements
formed towards the center. Not to scale.
PERCENT ABUNDANCE
While many elements have multiple
stable isotopes, that does not mean
that all stable isotopes will exist in
equal amounts in nature.
The percent abundance is a measure
of how much of a particular isotope
exists compared to the whole of that
element.
The reasons for each isotope’s
percent abundance vary from how
much is made in solar furnaces, the
formation of planets, planetary
geology, and others.
PERCENT ABUNDANCE
On Earth, most elements will have at least two decently-stable
isotopes that will both have fairly high abundance.
This is why when we want to know the mass of an element, we do not
use the mass of any one isotope – we use the average atomic mass,
a weighted average of the various possible isotopes.
For example, gallium exists as both Ga-69 (60.11%) and Ga-71
(39.89%). Therefore, if one wanted to know the mass of a piece of
gallium, it would be wrong to assume the mass to be that of Ga-69
(exactly 69 amu) or Ga-71 (exactly 71 amu). The mass will be in
between the two, though closer to 69, as it was the majority
component. The average atomic mass of gallium works out to be
69.723 amu.
PERCENT ABUNDANCE
Average atomic mass is can be calculated using the formula:
Where the fractional abundance is the percentage abundance of that
isotope expressed as a decimal.
This formula can be extended for as many isotopes are there are
present.
PERCENT ABUNDANCE
PERCENT ABUNDANCE
Example 2: Magnesium has 3 major isotopes: Mg-23 (23.99amu,
78.99%), Mg-24 (24.99 amu, 10.00%), and Mg-25 (25.98 amu,
11.01%). Find the Average Atomic Mass.
AMM = (0.7899 x 23.99 amu) + (0.1000 x 24.99 amu) + (0.1101 x 25.98 amu)
AMM = (18.9497 amu) + (2.499 amu) + (2.860398 amu)
AMM = 24.31 amu.
PERCENT ABUNDANCE
The atomic mass given on the
periodic table is the average
atomic mass in amu (atomic
mass units). Note that it is not
a whole number for most
elements as two or more
isotopes exist.
Many of the largest elements
are listed with a wholenumber AAM, often in
brackets, or have no AAM
listed at all. These are
elements that have only been
created artificially, and thus
have no “natural” abundance.
The number listed is of the
most-stable isotope created to
date.
IONS
An element is defined by its number of protons. Neutrons may vary,
but only contribute to the mass of the atom.
In an atom, the number of electrons is always equal to the
protons. This results in the positive and negative charges cancelling
each other out, and thus atoms always have a net charge of zero.
However, if the number of electrons and protons are different, the
particle is called an ion.
IONS
Atom (charge = 0)
Anion (charge < 0)
Cation (charge > 0)
IONS
The net charge of an ion is the summation of the positive and negative
charges (that is, the numbers of protons and electrons, respectively) in
an ion.
Net Charge = (protons) + (electrons)
 Note that electrons should be counted as negative in this formula.
Net charge is usually shown to the right of a chemical symbol as a
superscript. Ions are often shown encased in of square brackets. For
atoms (which have no charge) this is naturally omitted.
[F]- means a fluorine ion with a one minus charge. The “1” can be
omitted or included. Other integers (2, 3, etc) must be indicated, such
as in [O]2-.
IONS
To calculate the number of electrons in an ion, subtract the number of
protons (the atomic number, Z) from the net charge.
Electrons = Z – [net charge]
For example, a O2- ion:
 1) Oxygen has Z=8 (look up on periodic table)
 2) A 2- ion has a net charge of -2.
 Electrons = (-2) – (8)
 Electrons = (-10)
 Therefore, there are 10 (negatively-charged) electrons in a O2- ion.
SUMMARY
Protons:
 Protons are determined by the atomic number, Z
 If Z is not given, find using the mass number (pro = mass – neu), or from net charge
(pro = [net] – elec)
Neutrons:
 Neutrons are found by using the mass number and the atomic number
 Neutrons = (mass number) – (protons)
Electrons:
 Electrons are found by using the net charge and the number of protons
 Electrons = (net charge) – (protons)
PRACTICE
1)
How many electrons are in a Mg2+ ion?
2)
What is the net charge on a Boron-10 ion/atom that contains
2 electrons?
3)
What is the net charge on a Boron-10 ion/atom that contains
5 electrons?
4)
What element contains 2 electrons and a net charge of 2+?