Download PPT of Notes

Basic Atomic Structure –
History of the Atom
Early atomic structure
Greeks (400 – 300 B.C.)
Continuous theory of matter –
matter could be divided forever
without reaching a single smallest
indivisible unit
The discontinuous theory of
matter – there existed some smallet
piece of matter
John Dalton
John Dalton (1766 – 1844) “Father of Chemical
Atomic Theory”
Dalton was the first individual to really expound
upon the concept of the atom – this was a result of
evaluating the work of other scientists as well as
some of his own work involving the study of gases
Dalton’s theory published over the years 1803 –
1807 consisted of several main ideas:
Dalton’s Atomic Theory
1. Chemical elements are made up of atoms
– Elements are made up of tiny, discrete, indivisible and
indestructible particles called atoms.
• Derived from the Greeks
– These particles maintain their identity through chemical and
physical changes. They are only rearranged.
• Explained - Law of conservation of mass – developed by A.
Lavoisier
Note: Element – a chemical substance which cannot be further
simplified by chemical means
– a definition developed by A. Lavoisier
Dalton’s Atomic Theory cont’d
2. The atoms of an element are identical
- Atoms of the same element have the same properties, i.e.
weight
– Today, this concept has been proven incorrect with the
knowledge that elements have isotopes.
3. Atoms of different elements have different properties
- Atoms of different elements have different
properties including different weights.
– Again, this originated with the Greeks, however, it
was Dalton that determined what some the weight
values were.
Dalton’s Atomic Theory cont’d
4. Atoms combine in small whole number ratios
e.g. 1 to 1, 1 to 2, 2 to 3, etc.
– This coincides with Joseph Proust’s Law of Definite
Proportions which is also referred to as the Law of
Constant Composition.
– The law states that; compounds are composed of
elements combined in a fixed ratio by mass.
Dalton – other work
Dalton also developed another law known as the
Law of Multiple Proportions which states; that
the ratio of a fixed mass of one element when
combined with the multiple masses of another
element may be expressed in small whole
numbers.
In simpler terms, the same element can unite in
more than one way with another element to form
several compounds.
Dalton’s Model
John Dalton’s model of the atom
appeared as a spherical, indivisible and
indestructible unit
(~1807)
Electricity
Static electricity – is the build up of electrical
charge on an object (recognize most objects are
electrically neutral)
- being static implies that it is not moving
These objects of one charge are attracted to objects of the
opposite charge, however, objects with the same charge
repel each other
This relationship between charged objects is described by
the Law of Electrostatic Forces
Electrostatic force is the force observed between two
charged objects
Electricity cont’d
Electrically charged objects will attempt to return to
their neutral condition. Why?
This flow of electrical charge is referred to as
electrical current.
Electrical current is measured in units of
amperes
The ampere is defined as an electrical current
which transports one coulomb of charge past a
point in one second
The unit coulomb, C, represents a quantity
of charges equal to 6.24196 x 1018 charges
Electricity cont’d
The behavior of electrical charges is often
observed using electrodes.
Electrode – site or terminal of
electrical charge
Two types of electrodes; anode which is
positive and cathode which is negative
Coulomb’s Law
Fel = kQ1Q2
r2
where Fel is electrostatic force; k
is a constant with a value of 8.99 x 109Nm2/C2;
Q1 and Q2 are the quantity of charge on the two
objects; and r is the distance between the
charged objects
Understanding Coulomb’s Law
Implications:
- force is inversely proportion to the square of
the distance between the charges
- force is proportional to the product of the
magnitudes of the charges
- force is attractive for charges of opposite signs
and repulsive for charges of the same sign
Particles and Fields
A charged particle moving through a field (electrical or
magnetic) may have its path altered by the field
– It will be attracted toward the opposite pole and repelled from the
like pole
The particle may only be deflected if the field is
sufficiently weak, however, a strong field may actually
result in capture of the particle
It should be noted that deflection by a magnetic field is at
right angles (perpendicular) to the field while the
deflection by an electrical field is in the plane of that field
Deflection of Particles
Deflection of Nuclear Particles
(note charges)
Voltage
The strength of a field or current is related to the
potential energy of the electrons or quantity of
charge(s) present
This potential difference is measured in units
of volts – it is the “push” behind electron
movement
The volt is defined as a potential difference of 1
joule per coulomb
What is a joule?
Voltage
The greater the voltage, the greater the
potential energy of the system, and the
more work that may be done by that
system
Most household voltages are?
Techniques for working with small
moving charged objects
Experiment frequently require evacuated or
partially evacuated containers (metal or
glass) as gaseous molecules present may
interfere with the particle and its motion
Particle observation usually requires the
interaction/collision of the particle with
another substance
– use of a fluor or use of photographic film
Cathode Ray Tube (CRT)
OBSERVATIONS of the CRT
If an object is placed in the cathode ray a shadow will be cast upon the
wall at the glowing end of the tube
– Straight line of travel
CRT
The cathode ray can push a small paddle wheel up an incline against
the force of gravity
– The ray carried energy and could do work
The ray is deflected from straight line path by a magnetic field
Possible relationship between the two as the relationship between
magnetism and electricity hasn’t been discovered yet
CRT
Thomson Discovers the
Electron
J.J. Thomson deflects the cathode ray with an electrical field
– The rays bend toward the positive pole, therefore, they are negative
– 1897 announces that the corpuscles (electron) is a negatively charged
particle as it is deflected towards the positive plate in an electrical field
J.J. Thomson also Determined
Using a modified CRT Thomson worked on
determining the charge to mass ratio (e/m)
of the electron
The value obtained for the electron was:
e/m = 1.759 x 108 C/g
Charge to Mass Determination
Robert Millikan
Millikan performed the “oil drop” experiment in
1909 to determine the charge on the electron
Millikan Oil Drop
Millikan Oil Drop
**Millikan’s results showed that the value of ed
varied, but was always a multiple of a small
number.
- this resulted from the variation in
ionizations
**the small numerical multiple was
1.602 x 10-19 C
**Millikan believed this to be the smallest possible
charge, that belonging to the electron
Mass of the Electron
e___
e/m
=
= 1.602 x 10-19 C x
m
=
1.602 x 10-19 C
1.759 x 108 C/g
1 g______
1.759 x 108 C
9.107 x 10-28 g
Actual Electron Mass
Actual mass of the electron is
9.109 534 x 10-28 g
Memorize mass as 9.11 x 10-28 g
Proton
In 1886, Eugene Goldstein observed luminous
rays streaming through a cathode whose
modification was that it had holes
It also contained hydrogen gas at low pressure
The magnitude and direction of the rays in the
presence of a magnetic field indicated they were
positive, had a mass significantly larger than the
electron, and travel at a much lower velocity
Canal Rays
Proton
Credit for the discovery of the proton is
debatable pending upon source
- However, some people credit Rutherford with its
discovery
The proton’s mass was eventually identified
as 1.672 65 x 10-24 g
X-rays Discovered
Wilhelm C. Roentgen (1845 – 1923)
~1895 - Accidentally discovered x-rays when his
wife passed her hand in the path of x-rays
revealing her bone structure and wedding ring
This discovery inspired Becquerel to study
fluorescence and phosphorescence
The “Hand”
Radioactivity Discovered
A. Becquerel accidentally discovered radioactivity
in 1896 while working on fluorescence and
phosphorescence
Worked with potassium uranyl sulfate
– Radioactivity is defined as the spontaneous
emission of subatomic particles or energy by
disintegration of atomic nuclei
The Ore
Some types of radiation
Alpha (α) radiation – composed of helium nuclei traveling at fairly high
speeds
–
These particles carry a charge of +2 and have a mass of
~6.64 x 10-24 g
–
Alpha is a somewhat poor penetrating radiation
–
Equation symbol is 42He
–
Strong ionizing potential
Beta (β) radiation – composed of high speed electrons
–
These particles carry a charge of –1 and have a mass of 9.11
x 10-28 g
–
Beta has fair penetrating ability
–
Equation symbol is 0-1e
–
Good ionizing potential
Radiations cont’d
X-rays – are composed of electromagnetic radiation
–
–
They have good penetrating ability
Poor ionizing potential
Gamma rays – are similar to x-rays (electromagnetic radiation) but usually carry more
energy
–
–
–
They have very strong penetrating ability
Symbol is (γ)
Variable ionizing potential
Radiations
Gold Foil Experiment
Rutherford (1871 – 1973) along with
coworkers Geiger, Marsden, and Bohr
devised and performed the gold foil
experiment in 1911
The experiment:
A thin sheet of gold foil was enclosed by a
fluor covered circular shroud
– There was an opening in the shroud through
which alpha particles were shot at the gold foil
Gold Foil Experiment
Results of Gold Foil Experiment
1) The vast majority of alpha particles
passed through the foil without any
deflection from their path
2) Some of the alpha particles passed
through the foil with a minor deflection to
their path
3) A scant few (1 in 8,000) alpha particles
were deflected at large angles, in essence
bounced backwards
Rutherford’s Conclusions
1) Most of the alpha particles passed through the
foil without deflection because the vast majority
of the atom is empty space with a few electrons
in it
2) The few alpha particles deflected at minor
angles came close to a small bundle or core of
positive charge which were, therefore, repelled
3) The scant number of alpha particles being
bounced back was a result of an almost direct
collision with a very massive (extremely dense)
and positively charge region occupying a very
small volume
Atomic Models
• Dalton – spherical indivisible unit
•
Atomic Models
• Thomson – the plum pudding model
• a “sea” of positive charge with electrons scattered
throughout
• a divisible atom – subatomic particles existed
• an electrical nature associated with the atom
Atomic Models
Rutherford – planetary model
- the atom contained a small, positively charged, dense
core or center called the nucleus
- the electrons traveled around outside the nucleus
- most of the atom’s volume was actually empty space
- the nuclear diameter is about 10-4 the diameter of the
atom
Problems w/ planetary model
• Criticisms of the planetary model of the atom
• Why weren’t the electrons pulled into the positively
charged nucleus of the atom?
– Rutherford responded by stating the electron’s motion prevents it
from being pulled into the nucleus much the same as the planets
aren’t pulled into the sun or the moon into the earth
• According to classical mechanics (physics) charged
particles moving in a curved path should emit energy
(light) or some other form of electromagnetic radiation.
Eventually, they would lose enough energy to be pulled
into the nucleus
Neutron
In 1928, a German physicist, Walter Bothe, and his student, Herbert
Becker, took the initial step in the search for the neutron. They
bombarded beryllium with alpha particles emitted from polonium and
found that it gave off a penetrating, electrically neutral radiation,
which they interpreted to be high-energy gamma photons.
Neutron cont’d
• Eventually, the neutron was discovered in
1932 when James Chadwick used
scattering data to calculate the mass of
this neutral particle.
– Chadwick is credited with his discovery
Chadwick’s Experiment
Basic Structure of Modern Atom
There are three major subatomic particles:
The proton with a stadardized charge of +1, a mass 1.673
x 10-24 g, and located in the nucleus
The neutron with a charge of 0, a mass of 1.678 x 10-24 g,
and located in the nucleus
The neutron is slightly larger than the proton
By what?
Major Subatomic Particles
cont’d
The electron with a standardized charge of
–1, a mass of 9.11 x 10-28 g, and located
outside the nucleus in what is called the
electron cloud
The electron is ~ 1 / 1837 the mass of
the proton
Facts about the Basic Structure
of Modern Atom
Atoms are usually found in their neutral
state, having equal numbers of protons
and electrons
# protons = # electrons for NEUTRAL
atoms
Most of the atom is composed of empty
space (vacuum), the volume of the atom
being established by the electron cloud
Facts about the Basic Structure
of Modern Atom
The nucleus is an extremely small, extremely
dense, positive core of the atom which contains
the vast majority of the atom’s mass
Nuclear density is about 1013 to 1014
g/cm3
Why don’t the electrons follow Newton’s First law
and fly off and out of the atom?
Electrostatic force - Coulombic
attraction
Facts about the Basic Structure
of Modern Atom
Atomic diameters are expressed on the order of angstroms
1 angstrom = 10-10 m
Symbol for angstrom is Å
Atomic nuclei have diameters that are on the order of 10-4
angstroms
The identity of an element is established by the number of
protons contained in the nuclei of its atoms
This number of protons is referred to as the atomic
number of the element
The symbol for the atomic number is Z
Facts about the Basic Structure
of Modern Atom
Atoms may lose or gain electrons without
any change in element identity, BUT any
change in number of protons will create a
new identity for that atom
– Loss of electrons results in ions with a
positive charge (called cations)
– Gain of electrons results in ions of a negative
charge (called anions)
Isotopes
Isotopes are atoms of the same element which differ in
mass
This mass difference is due to possession of differing
numbers of neutrons
Recognize, nuclear charge cannot change
Isotopes of an element have different mass numbers
Mass number (symbol A) is equal to the sum of major
nucleons
A nucleon refers to any particle in the nucleus of an
atom
Mass Number
Therefore, mass number equals the sum
of protons and neutrons
Mass number = Z + #n or
A = Z + #n
since A is the symbol for the mass
number
– To find the number of neutrons in an
isotope use;
#n = A - Z
Representing Isotopes
Isotopes have the general form of; AZX, where X is the
chemical symbol of the element
e.g.
1H;
14C;
206Pb
often the value of
Z isn’t expressed – why
not?
If the isotope is an ion a charge with a + or – sign would be
in the right superscript
e.g. 5726Fe+2
Naming Isotopes
The term nuclide is used to refer to a specific atom of an
isotope
– an atom with specific number of protons and neutrons
Hydrogen has three isotopes with special names:
1H - protium;
2H - deuterium;
3H
- tritium
Other naming involves the element name, a dash, and the mass
number
– e.g. 57Fe would be iron – 57 (implies neutral)
– e.g. 57Fe+2 would be the iron – 57 +2 cation (or ion)
– E.g. 33S-2 would be the sulfur – 33 -2 anion (or ion)
Atomic Mass Scale
Atoms are too small too measure their mass directly,
however, their relative masses may be determined.
Compounds are evaluated for the combining ratios of the
atoms. After which upon decomposition, an analysis of the
mass of each element may be established.
From this information, the relative mass of each element
may be determined.
Atomic Mass Scale
In order to provide uniformity, the atomic
mass scale was established.
The current reference standard is the
isotope carbon – 12 (12C). It is assigned a
mass of exactly 12 atomic mass units
(a.m.u. or u.).
1 nuclide 12C = 12.000 000 000… a.m.u.
What is an Atomic Mass Unit?
The atomic mass unit is the unit of the
atomic mass scale. It is defined as 1/12th the
mass of the carbon – 12 nuclide.
It is abbreviated as a.m.u. or u.
1 a.m.u. = 1/12
12C
nuclide
Average Atomic Mass
In looking at the masses of the elements on
the periodic table it is evident that most
values listed aren’t close to being whole
numbers
Why?
These values are actually averages that
represent all the naturally occurring
isotopes and the relative abundance in
which they are found in nature
Average Atomic Mass
The average atomic mass may be
determined by calculation from any
analysis yielding the actual masses of the
isotopes and some relative proportion of
their presence (an actual number of
nuclides or as percentages of nuclides)
A “rough estimate” may be obtained if the
mass numbers are utilized
Average Atomic Mass Equations
Ave. At. Mass = [(% x isotope mass) + (% x isotope mass) + …..]
Total %
Ave. At. Mass = [(#nuclides x isotope mass) + (#nuclides x isotope mass) + ..]
Total # nuclides
Mass Spec
The mass spectrometer or mass spectrograph is a useful instrument
in determining chemical analysis or analysis of isotopes
The atoms of material are ionized giving them a positive charge
Their charge is usually +1 but may sometimes be greater
The ions travel through the magnetic field assuming a curved path
which directs them onto a detector or film plate
Number of hits at a particular spot gives the abundance
(concentration)
**Placement of hits leads to the calculation of mass to charge
and eventually identity
Mass Spec diagram
Deflection of particles
Mass Spec
Different ions are deflected by the magnetic field by different
amounts. The amount of deflection depends on:
– the mass of the ion. Lighter ions are deflected more than
heavier ones.
– the charge on the ion. Ions with 2 (or more) positive
charges are deflected more than ones with only 1 positive
charge.
These two factors are combined into the mass/charge ratio.
Mass/charge ratio is given the symbol m/z (or sometimes
m/e).
– For example, if an ion had a mass of 28 and a charge of
1+, its mass/charge ratio would be 28. An ion with a mass
of 56 and a charge of 2+ would also have a mass/charge
ratio of 28.
Mass Spec
In the diagram above, ion stream A is most deflected - it will
contain ions with the smallest mass/charge ratio. Ion
stream C is the least deflected - it contains ions with the
greatest mass/charge ratio. It makes it simpler to talk
about this if we assume that the charge on all the ions is
1+. Most of the ions passing through the mass
spectrometer will have a charge of 1+, so that the
mass/charge ratio will be the same as the mass of the ion
Note: You must be aware of the possibility of 2+ (etc) ions,
but the vast majority of first level questions will give you
mass spectra which only involve 1+ ions. Unless there is
some hint in the question, you can reasonably assume that
the ions you are talking about will have a charge of 1+.
Mass Spec
Assuming 1+ ions, stream A has the lightest
ions, stream B the next lightest and stream
C the heaviest. Lighter ions are going to
be more deflected than heavy ones.