Atoms and Elements
Instructor: Dr. Niloufar Choubdar
• About 2,500 years ago, Democritus was one of the first early Greek philosophers to propose
"atoms“. He suggested that atomos, or atomon—tiny, indivisible, solid objects—make up all
matter in the universe.
• Democritus then reasoned that changes occur when the many atomos in an object were
reconnected or recombined in different ways. Democritus even extended this theory, suggesting
that there were different varieties of atomos with different shapes, sizes, and masses.
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• The modern atomic theory, proposed about 1803 by the English chemist John Dalton, is
a fundamental concept that states that all elements are composed of atoms:
• All matter is composed of extremely small particles called atoms.
• Atoms of a given element are identical in size, mass, and other properties. Atoms of
different elements differ in size, mass, and other properties.
• Atoms cannot be subdivided, created, or destroyed.
• Atoms of different elements can combine in simple whole number ratios to form
chemical compounds.
• In chemical reactions, atoms are combined, separated, or rearranged.
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• Dalton's atomic theory has been largely accepted by the scientific community, with the exception
• of three changes:
• (1) an atom can be further subdivided
• (2) all atoms of an element are not identical in mass
• (3) using nuclear techniques, we can create or destroy atoms by changing them into other atoms.
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• for almost 100 years, it seemed as if Dalton's Atomic Theory was the whole truth.
• However, in 1897, J. J. Thomson conducted some research that suggested the small, negatively
charged particles making up the cathode ray were pieces of atoms. Today we know them as
electrons.
• The existence of protons was also known, due to the fact that atoms were neutral in charge.
• The next step was to figure out how these particles were arranged in the atom.
• J.J. Thomson developed what became known as the"plum pudding" model in 1904.
• In such model, the electrons were embedded in a uniform sphere of positive charge
like plum stuck into pudding.
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Rutherford experiment
• Ernest Rutherford (1871‐1937) about five years later developed a new model.
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Observations from Rutherford experiment
• According to the plum pudding atomic model, in which an atom's mass and charge are uniformly
distributed throughout the atom, the scientists expected that all of the alpha particles would pass
through the gold foil with only a slight deflection or none at all.
• Most of the alpha particles were indeed not deflected, a very small percentage (about 1 in 8000
particles) bounced off the gold foil at very large angles.
• Some were even redirected back toward the source.
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Rutherford model
• Nucleus is the tiny, dense, central core of the atom and is composed of protons and neutrons, and
electrons are distributed around the nucleus in the atom.
• Because most of the atom was empty, so the vast majority of the alpha particles had passed
through.
• For small number of alpha particles (positively charged) that bombed into the nucleus, they were
highly deflected because of the powerful repulsion force of the nucleus with positive charge as well.
• Comparing the size: If we could blow up an atom to be the size of a large professional football
stadium, the nucleus would be about the size of a marble.
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Nuclear Atomic Model
• The nucleus (plural, nuclei) is a positively charged region at the center of the atom. It consists of
two types of particles which have a positive electric charge, and neutrons, which are neutral in
electric charge. Outside of the nucleus, an atom is mostly empty space, with orbiting negative
particles called electrons.
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Comparing size of nucleus
• The nucleus of the atom is extremely small. Its radius is only about 1/100,000 of the total radius
of the atom.
• Electrons have virtually no mass, but protons and neutrons have a lot of mass for their size.
• The nucleus has virtually all the mass of an atom. Given its great mass and tiny size, the nucleus is
very dense.
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• Protons
• Protons are found in the nucleus.
• Protons have a positive electrical charge of one (+1) and a mass of 1.007 atomic mass unit (amu),
• Neutrons
• Neutron has no charge. The mass of a neutron is slightly greater than the mass of a proton, which
is 1.0087 atomic mass unit (amu).
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Electrons
• Electrons are extremely small. The mass of an electron is only about 1/2000 the mass of a proton.
• Electrons have an electric charge of −1, which is equal but opposite to the charge of a proton
• All atoms have the same number of electrons as protons.
• electrons are found outside the nucleus.
• Particle
Mass (amu)
Relative Charge
• proton
1.0073
+1
• neutron
1.0087
0
• electron
0.00055
-1
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Atomic number
• The number of protons in an atom is called its atomic number (Z). This number is unique for
atoms of a given element.
• An element's atomic number also tells you how many electrons are in a neutral atom of that
element.
• For example: hydrogen has an atomic number of 1. This means that an atom of hydrogen has one
proton, and one electron
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mass number (A)
• The mass number (A) of an atom is the total number of protons and neutrons in its nucleus.
• The mass of the atom is also in the unit of atomic mass unit (amu).
• Because electrons have virtually no mass, just about all the mass of an atom comes from protons
and neutrons.
• mass number A = number of protons + number of neutrons
• Example:
• Helium atoms have 2 neutrons in addition to 2 protons:
• Therefore, the mass of most helium atoms is 4.
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Isotopes
• Because the number of neutrons can vary for a given element, atoms with the same number of
protons but different numbers of neutrons are called isotopes.
• Atomic symbol is the symbol to specify all the above information for an element together:
• X: the symbol of the element
• A: the mass number (number of protons plus neutrons)
• Z: the atomic number (number of protons)
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• For example, for helium:
• one isotope of helium element has 2 neutrons. The other isotope of helium element with three
neutrons has the symbol of He
• Another common example of isotope is for Hydrogen:
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Atomic mass
• The average atomic mass of an element is:
• The weighted average mass of the atoms in a naturally occurring sample of the element.
• The atomic mass listed in the periodic table is the average atomic mass, with unit of atomic mass unit
(amu).
• Calculating Average Atomic Mass
• The exact mass of a certain naturally occurring isotope, and the relative abundance (the fraction of
• a given isotope of an element) of that isotope.
• The average atomic mass can be calculated by the following equation:
• Average Atomic Mass = (fraction of isotope 1)(mass of isotope 1)+( fraction of isotope 2)(mass of
• isotope 2) + fraction of isotope 3)(mass of isotope 3) + ……
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• Example
• Magnesium has three naturally occurring isotopes. Mg‐24 with mass 23.99amu and a natural
• abundance of 78.99%, Mg‐25 has a mass of 24.99amu and natural abundance 10.00%, and Mg‐
• 26 with mass of 25.98amu and a natural abundance of 11.01%. What is the average atomic mass
• of Magnesium?
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• Example Copper has two naturally occurring isotopes, Cu‐63 with mass 62.9396 amu and Cu‐65
• with mass 64.9278 amu. The average atomic mass of copper is 63.55 amu, what is the natural
• abundance of each isotope for copper?
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