ADVANCED CHEMISTRY
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Know the history of atom and how it was dicovery.
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Be able to understand the experiment that took place in the past.
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To be able to explain what is atomic number, mass number, and isotopes.
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Fifth century BC Greek philosopher Democritus believed that matter was
comprised of minimum particles which could not be subdivided and
maintain the character of the element, and called them atomos. The
theory that atoms are the fundamental building blocks of matter
reemerged in 1808, championed by John Dalton in his atomic theory.
1.
Each element is composed of extremely small particles called atoms.
2.
All atoms of a given element are identical to one another in mass and
other properties, but the atoms of one element are different from the
atoms of all other elements.
3.
Atoms of an element are not changed into atoms of a different
element by chemical reactions; atoms are neither created nor
destroyed in chemical reactions.
4.
Compounds are formed when atoms of more than one element
combine; a given compound always has the same relative number and
kind of atoms.
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Law of Definite Proportions
 States that a compound is always composed of the same elements at
fixed proportions by mass.
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Law of Multiple Proportions
 States that when two elements combine to form more than one
compound, the mass of one element.
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Law of the Conservation of Mass
 States that the total mass of substances present at the end of a
chemical process is the same as the mass of substances present
before the process took place.
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In the 1890’s there was much interest in the study of radiation.
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Cathode rays, were generated in an evacuated tube. Streams of these
negatively charged particles were found to emanate from cathode tubes,
causing fluorescence.
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They were emitted by the cathode and passed through a hole in the anode,
and could be deflected by either a magnetic field and/or an eletrostatic field.
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J. J. Thomson is credited with their discovery (1897) and he determined the
charge/mass ratio of these rays to be –1.76 x 108 C/g.
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In his Millikan Oil-Drop Experiment, Robert Millikan (University of Chicago)
in 1909 used charged oil droplets and an electrostatic field to determine the
charge of an electron to be –1.6022 x 10-19 C. Hence was able to calculate the
mass of an electron to be 9.10 x 10-28g
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Radioactivity is the spontaneous emission of radiation by an atom.
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It was first observed by Henri Becquerel and Marie and Pierre Curie also
studied it.
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In 1895 Röntgen discovered rays, emitted by materials activated by
cathode rays that were uncharged.
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Three types of radiation were discovered by Ernest Rutherford: α
particles, β particles and γ rays.
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In 1910 Rutherford found that alpha-particles passed through gold foil
undeflected and occasionally slightly deflected, but sometimes
completely bounced back.
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He formulated a model where the atom was comprised of mainly empty
space, but with a concentrated positive nucleus.
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Protons were discovered by Rutherford in 1919.
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Protons, the positively charged particles contained in the nucleus were
later determined to have the same magnitude of charge as an electron,
but has a mass of 1.67262 x 10-24g.
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Mass discrepancies had Rutherford suppose the presence of other massy
particles in the nucleus.
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James Chadwick in 1932 discovered the neutron which has a mass
slightly bigger than the proton, but without charge (1.67493 x 10-24g).
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Atomic number (Z) is the number of protons in its nucleus
which equals its number of electrons if it is uncharged.
 The chemical properties of an element are determined by its atomic number,
and the mass seldom plays a role in the chemistry of an element.
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Mass number (A) is the sum of the protons and neutrons in
an atom.
 The mass of an atom in atomic mass units (amu) is the total number of
protons and neutrons in the atom.
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Atoms with the same number of protons (atomic number)
but different number of neutrons (nucleon or mass number)
are called isotopes.
 Most elements are comprised of several natural isotopes.
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Atomic and molecular masses can be measured with great accuracy
using a mass spectrometer. Because in the real world we use large
amounts of atoms and molecules, we use average masses in calculations.
Average mass is calculated from the isotopes of an element weighted by
their relative abundances.
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1. Lecture: Theodore E. B., Eugene, H. L. H., Bruce E. B., Catherine M., Patrick W.,
(2011). Chemistry: The Central Science (12 Ed). Prentice Hall. USA.
2. Laboratory: Theodore E. B., John H. N., Kenneth C. K., Matthew S. (2011).
Laboratory Experiments for Chemistry: The Central Science (12 Ed). Prentice Hall.
USA.
3. Theodore E. B., (2011). Solutions to Exercises for Chemistry: The Central
Science. Prentice Hall. USA.
4. John M., Robert C. F. (2010). Chemistry (4 Ed): Prentice Hall Companion
Website.
http://wps.prenhall.com/esm_mcmurry_chemistry_4/9/2408/616516.cw/index.ht
ml
5. Chemistry Online at
http://preparatorychemistry.com/Bishop_Chemistry_First.htm
6. Chemistry and You at
http://www.saskschools.ca/curr_content/science9/chemistry/index.html
7. Teachers Notes
8.www.study.com