ATOMIC
STRUCTURE
History of the Development of
Modern Atomic Theory
·
Ancient Greeks believed that matter is
composed of a limited number of basic, fundamental component substances called
elements: earth, wind, water and fire.
·
Democritus (460-370) states that if
matter is cut into consecutively smaller pieces, at some point it can be cut no
longer and still retain its particular properties. Such a tiny particle was
called an atom (ancient Greek for
“uncuttable”). Democritus’ theory was philosophically correct, but not
generally accepted (matter was viewed to be continuous): there were no
experiments to test the theory.
·
John
Dalton’s Atomic Theory
After
considerable experimentation into the manner in which the chemical composition
change in a definite way during chemical reactions, John Dalton, an English
School teacher, proposed a theory involving the concept of indestructible
atoms.
Dalton’s
Atomic Theory has 5 postulates:
1. Matter
consists of definite particles called atoms.
2. All
atoms of one particular element are identical in size, mass and chemical
properties.
3. Atoms
of different elements have different masses: this being the distinguishing
feature.
4. Compounds
are formed by binding together atoms of different elements in definite,
whole-number ratios of atoms.
5. All
atoms are indestructible. Compounds do not result in atom’s creation or
destruction, but involves only the separation, combination, or rearrangement of
atoms.
These
postulates were considered as the basic concepts for more investigation and
studies done by most of the scientist and as a result, they found out that some
of these postulates were proven correct and others were disproved. Dalton’s
atomic theory and the series of developments led to the modern atomic theory of
atomic structure.
·
The
Cathode Ray Tube
The
cathode ray tube was one of the devices utilized in the discovery of cathode
rays, which was first studied by J. Plucker in 1858. These rays have properties
of traveling in a straight line away from the cathode (negative electrode) to
the anode (positive electrode), therefore the rays must be negatively charged.
J.J. Thomson further studied these negatively charged particles called
electrons by studying the degree of deflections of the cathode rays in
different magnetic and electric fields and found the charge to mass ratio of
electrons to be -1.75882 x 108 coulombs per gram. This ratio is the
same to all nature of the cathode rays (regardless of the type of the gas)
which indicates that electrons are fundamental particles found in all matter.
Between
1908 and 1917 R.A. Millikan succeeded in measuring the charge of the electron
through his oil-drop. From the mass and total charge of the oil drop, Millikan
was able to determine the charge of the electron – 1.602 x 10-19 C.
he was also able to calculate the mass of the electron: 9.109 x 10-28g.
Mass of an electron - -1.6022 x 10-19C - 9.10 x 10-28g
-1.76 x 108C/g
·
Thomson’s
Plum Pudding Model
The
atom was, on the whole, electrically neutral. Thus, there must also be a source
of positive charge within the atom, which exactly cancels out the negative
charge of the electron.
Another
English chemist, Joseph John Thomson proposed that the atom might be thought of
as a tiny sphere throughout which positive charge is evenly deposited, with
just enough negatively charged electrons randomly scattered about to balance
the positive charge.
·
Rutherford’s
“Nuclear Atom”
In
1911, English Physicist Ernest Rutherford and co-workers carried out an experiment,
which destroyed the “Plum Pudding Model” and set the stage for the modern
picture of the atom.
While
studying alpha particles (a form of radiation, positively charged particles
with a mass 7,300 x that of an electron), Rutherford found that these
relatively massive particles could be deflected by collision with other equally
massive particles.
Shooting
a beam of alpha particles at a very thin gold foil, most of the alpha particles
passed through unaffected, but some had their path or trajectory greatly
changed, some being reflected backward.
Rutherford’s Interpretation of Results
1.
The
large deflections of the alpha particles could only be explained by a massive
center of positive charge that would repel the positively charged alpha
particles. The deflected alpha particles come close to this center of charge
and its path is bent, whereas the reflected alpha particles score a “direct
hit” on this center and bounces toward the alpha emitter.
2.
Most of the alpha particles pass directly
through the foil because most of the atom is “empty space”. The nuclear atom
has a heavy, positively charged nucleus (Greek for kernel) that occupies only a
small volume relative to the atom as a whole. The nucleus is 1/100,000 the size
of the atom.
3. The
nucleus must contain protons, a positively charged particle to balance the
charge of electrons. The charge of the proton is +1; the charge of the electron
is -1.
4. The
atom contains as many protons in the nucleus as electrons.
Nucleus
– the positively charge central part of the atom. It is very small with a
diameter of 10-3 Aº
(10-6 nm). Although it is very small compared to the whole atom, it
contains about 99% of the whole mas of the atom. Its density is 1.0 x 1014g/m3. Inside are the protons the neutrons
(collectively nucleons) surrounded by diffuse electrons.
Discovery
of Proton
In 1886, Eugene Goldstein found that
cathode ray tube also a stream of positively charged particles that moves
toward the cathode. These rays were called positive rays or canal rays. Studies
showed that different elements gave the same positive rays and this led to an
idea that it is a fundamental particle and that there is a unit of positive
charge called proton. The change is equal in magnitude but opposite in sign to
the charge of electrons.
Discovery
of Neutrons
In
1932, J. Chadwick identified the third fundamental particles of the atom as
neutron a particle which has about the same mass as the proton but without
electrical charge. The neutron, like the proton is found in the nucleus of the
atoms.
Particles
of Atoms
Protons – is
a positively charged particle, with a mass of 1.672 x 10-24gram or
1.0073 AMU. It is identical to the nucleus of the hydrogen atom and is 1,836
times heavier than an electron. Although heavier than an electron, it is much
smaller in size about 2/5 as big.
Electrons –
are negatively charged particles with a mass of 9.10 x 10-28 gram.
It has a charge of -1.602 x 10-19 coulombs, the smallest quantity of
electrical charge ever observed. All electrons have the same charge. It
balances exactly the positive charge of proton.
Neutrons – are
particles with no electrical charges and each has a mass of 1.675 x 10-24 grams
or 1.0087 AMU which is about the same mass as the proton. Because it is
electrically neutral, electrical forces of its atomic mass is 6.0 not repel it,
which makes it an ideal projectile for bombarding atoms to initiate nuclear
reactions.
Other particles were detected by physicists in their
experiment and some of these particles with mass the same as that of the
electron and meson particles, 200 times heavier than an electron with a
positive and negative charges. These particles were found to be unstable.
Summary
1. Atoms
consist of one or more sub-atomic particles called protons, neutrons, and electrons.
2. The
mass of an atom is concentrated in the tiny positively charged center core
called the nucleus.
3. The
nucleus of an atom contains one or more protons (=1 charge) and one or more
neutrons (no charge).
4. The
total positive charge of the nucleus is equal to the number of protons found in
the nucleus.
5. One
or more electrons (-1 charge) are in constant motion outside the nucleus.
6. The
number of electrons outside the nucleus is equal to the number of protons found
within the nucleus so that the atom is neutral overall.
7. All
atoms of the same element have the same number of protons in the nucleus and
the same number of electrons moving around the nucleus.
8. An
atom is mostly empty space.
9. Atoms
of the 110+ known elements differ in the number of protons, neutrons, and
electrons they contain.
Atomic
Number, Z
Atomic number is equal to the number of protons that an
atom contains. It indicates the nuclear charge of that atom of +1 because it
has only one proton. The atomic number of helium is 2 and the nuclear charge is
+2 because it has 2 protons. Uranium has atomic number of 92 and a nuclear
charge of +92 because it has 92 protons.
Atomic Mass or Mass
Number, A
This number is equal to the sum of the number of protons
and the number of neutrons present in the nucleus of an atom of an element. The
mass of H is 1 because it has one proton and no neutron. The atomic mass of He
is 4 because it has 2 protons and 2 neutrons. Lithium has 3 protons and 3
neutrons; therefore its atomic mass is 6.
Mass
Number – number of protons + number of neutrons
-
Atomic number + number of neutron
To
denote the atomic number and mass number of an atom of an element X is as
follows:
AZX
Where the superscript A
is the mass number and the subscript Z is the atomic number.
Example: 2311Na
Because of the very small values for
the mass of atoms, grams are an inconvenient unit of measurement. The standard
unit for measuring the mass of atoms is the atomic mass unit or amu.
1
amu – 1/12 the mass of a 12C atom – 1.66 x 10-24 g
Mass
of one 12C atom – 12 amu (exact by definition)
This is very close to the mass of a
proton and a neutron. Thus the mass of an atom, in amu, is very close to the
mass number.
Atomic mass is usually not whole
number. The reason is that atomic mass values are average atomic masses of the
different isotopes of an element taken in the proportion of their abundance.
Thus, if the atomic mass of an element is with decimal point it means that the
element has isotopes. Isotopes are atoms that have the same atomic number but different
mass.
Example:
Chlorine has isotopes Cl-35, which has about 75% and Cl37,
which is 25% of all atoms present in the world.
To compute the atomic mass: 0.75 x 35 amu = 26.25
0.25
x 37 amu = 9.25
35.5 atomic mass of Cl
Practice
Exercise: A certain element consists of the following
isotopes. Determine the average atomic mass and identify the element.
Mass (amu) %
Abundance
203.973 1.40
205.9745 24.10
206.9759 22.10
207.9766 52.40
RECORDED LECTURES
How to Calculate Atomic Mass Practice Problems
https://www.youtube.com/watch?v=ULRsJYhQmlo
Uses of radioactive isotopes - Chemistry
https://www.youtube.com/watch?v=E4B94zCY4ok
RECORDED LECTURES
Walang komento:
Mag-post ng isang Komento