Download Review Notes for Atomic Structure and Radioactivity Test on Friday

Review Notes for Atoms- The Building Blocks of Matter Test
1. Atoms are the smallest particles of an element that retain the chemical identity (properties)
of that element.
Early Models of the Atom:
2. The ancient Greek philosophers, Aristotle and Plato described matter as continuous and
without a definite unit. Democritus claimed that matter is discontinuous and if you keep
dividing it you will reach a final definite unit that gives matter its properties = the atomos.
3. Dalton’s atomic theory of matter forms the basis of our understanding of chemistry. It was
based on the law of conservation of mass – stated by Antoine Lavoisier in 1789, the law of
definite proportions – stated by Joseph Proust in 1799, and the law of multiple proportions –
Stated by John Dalton in 1803.
Discovering Atomic Structure:
4. Atoms contain particles that have electrical charge. Particles that have different charges
attract each other. Particles that have similar charges repel each other.
5. Based on his experiments that looked at the effect of electric and magnetic fields on
cathode rays, Thomson concluded that a cathode ray consist of negatively charged particles,
which he named electrons. Thomson’s model of the atom is described as “ The Plum Pudding
Model” or “The Chocolate Chip Cookie Model”.
6. Bacquerel discovered radioactivity, the spontaneous emission of particles and/or energy
from an element. Radioactivity is the product of an imbalance between the number of
protons and neutrons in the nucleus and results in fundamental changes in the nucleus of the
atom.
7. Based on his experiment of scattering alpha particles by gold foil, Rutherford described the
atom as an entity having most of its mass in a tiny, positively charged nucleus at the center
with electrons moving in an empty space surrounding the nucleus.
Modern Atomic Theory:
8. Atoms are composed of subatomic particles known as protons, neutrons , and electrons. The
nucleus is in the center of the atom. It contains protons and neutrons - also known as
nucleons.
9. The electrons are around the nucleus forming an electron cloud.
10. Most of the atom is an empty space (electron cloud) - was found in Rutherford's gold foil
experiment.
CHARGE
MASS
PROTON
positive
~1 atomic mass
unit (a.m.u.)
NEUTRON
none
~1 atomic mass
unit (a.m.u.)
LOCATION
In the Nucleus
In the Nucleus
ELECTRON
negative
hardly any;
1/1837 atomic
mass unit
Considered 0
In the Space
Surrounding the
Nucleus
11. One atomic mass unit is equal to 1/12 the mass of isotope carbon-12 standard used.
12
C which is the
12. If you want to know the symbols for proton, neutron, and electron, look at TABLE O. Table
O shows the symbols, charge and mass for all subatomic particles - the bottom number in
front of the symbol is the charge and the top number is the mass in a.m.u.
13. Atomic number - the number on the bottom before the element’s symbol - is unique to
each element and is equal to the number of protons, which is equal to the number of
electrons in a neutral atom.
14. Mass number - equal to the total number of protons plus neutrons, and appears at the top
before the symbol or in some cases after the symbol, Ex: C-12 or Na-23.
REMEMBER: To find number of neutrons:
Number of Neutrons = Mass number - Atomic number
15. An ion is an atom or a group of atoms with a net electrical charge. An atom becomes an ion
when it loses or gains electrons. An atom that loses electrons – has less electrons than
protons and therefore, it has an overall positive charge. An atom that gains electrons – has
more electrons than protons and therefore, it has an overall negative charge.
16. Isotopes – These are atoms of the same element with the same number of protons but
different number of neutrons. Therefore, the atomic number (bottom) is the same and the
13
14
mass number (top) is different, Ex: 12
6 C and 6 C and 6 C are all different isotopes of the
element Carbon .
17. Atomic mass - the weighted average mass of all the naturally occurring isotopes of the
element. The average is weighted according to the percent abundance of each one of the
isotopes. To calculate atomic mass we need the atomic mass of each of the isotopes and
their percentage.
Atomic Mass = Atomic Mass (A) x % abundance(A) + Atomic Mass (B) x % abundance(B)
EXAMPLE: An element has 2 isotopes - 75% of X-35 with a mass of 35.092 amu and 25%
of X-37 with a mass of 36.948 amu. What is the atomic mass? What is element X?
75 x 35.092amu 25 x 36.948amu

 35.556 amu
100
100
To find element X, look on the periodic table for the element with this atomic mass = Cl.
Radioactivity:
18. Radioactivity is particle or energy emission due to nuclear disintegration. You will find the
symbol, charge and mass for the different particles on Reference Table O. Radioactive
decay is spontaneous and cannot be controlled.
19. Many elements have at least one naturally occurring radioactive isotopes, or radioisotopes.
Elements with atomic number higher than 83 have no stable isotopes.
20. Half-life of a radioisotope is the time it takes for one half of a sample of that isotope to
decay. Radioisotope will never decay to an amount of 0 , it will simply come closer and closer
to that amount. The half-life for each radioisotope is constant; it cannot be changed by
changes in
pressure, temperature, or chemical combination. It depends only on the
nature of the radioisotope. The Half-Life of selected Radioisotopes is Listed in
Reference Table N. The formula for Radioactive Decay, you will find in Reference
Table T – (Formulas’ table).