CHM 101 LECTURE NOTES
Course Synopsis:
1. Modern electronic theory of atoms
2. Radioactivity
3. Chemical bonding
MODERN ELECTRONIC THEORY OF ATOMS
In our study of atomic structure, we look first at the fundamental particles. These are the basic building
blocks of all atoms. Atoms, and hence all matter, consist principally of three fundamental particles:
electrons, protons, and neutrons. Knowledge of the nature and functions of these particles is essential to
understanding chemical interactions. The relative masses and charges of the three fundamental particles are
shown in Table
Fundamental Particles of Matter
The mass of an electron is very small compared with the mass of either a proton or a neutron. The charge
on a proton is equal in magnitude, but opposite in sign, to the charge on an electron. The proton is a
fundamental particle with a charge equal in magnitude but opposite in sign to the charge on the electron.
Its mass is almost 1836 times that of the electron.
Note: Many other subatomic particles, such as quarks, positrons, neutrinos, pions, and muons,
have also been discovered.
Rutherford was able to determine the magnitudes of the positive charges on the atomic nuclei. The picture
of atomic structure that he developed is called the Rutherford model of the atom.
Atoms consist of very small, very dense positively charged nuclei surrounded by clouds of
electrons at relatively great distances from the nuclei.
ATOMIC NUMBER
Moseley (1887–1915) showed that the X-ray wavelengths could be better correlated with the atomic
number. On the basis of his mathematical analysis of these X-ray data, he concluded that
each element differs from the preceding element by having one more positive charge in its nucleus.
We now know that every nucleus contains an integral number of protons exactly equal to the number of
electrons in a neutral atom of the element. Every hydrogen atom contains one proton, every helium atom
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,contains two protons, and every lithium atom contains three protons. The number of protons in the nucleus
of an atom determines its identity; this number is known as the atomic number of that element.
MASS NUMBER AND ISOTOPES
Most elements consist of atoms of different masses, called isotopes. The isotopes of a given element contain
the same number of protons (and also the same number of electrons) because they are atoms of the same
element. They differ in mass because they contain different numbers of neutrons in their nuclei.
Isotopes are atoms of the same element with different masses; they are atoms containing the same
number of protons but different numbers of neutrons.
The mass number of an atom is the sum of the number of protons and the number of neutrons in its
nucleus
Mass number = number of protons + number of neutrons
= atomic number + neutron number
The composition of a nucleus is indicated by its nuclide symbol. This consists of the symbol for the element
(E), with the atomic number (Z) written as a subscript at the lower left and the mass number (A) as a
superscript at the upper left, AZE.
Note of Emphasis
✓ The atomic number, Z, is an integer equal to the number of protons in the nucleus of an atom of the
element. It is also equal to the number of electrons in a neutral atom. It is the same for all atoms of
an element.
✓ The mass number, A, is an integer equal to the sum of the number of protons and the number of
neutrons in the nucleus of an atom of a particular isotope of an element. It is different for different
isotopes of the same element.
✓ Many elements occur in nature as mixtures of isotopes. The atomic weight of such an element is the
weighted average of the masses of its isotopes. Atomic weights are fractional numbers, not integers.
Example:
1. Determine the number of protons, neutrons, and electrons in each of the following species. Are the
members within each pair isotopes?
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, Calculation of Atomic Weight
Three isotopes of magnesium occur in nature. Their abundances and masses, determined by mass
spectrometry, are listed in the following table below. Use this information to calculate the atomic weight
of magnesium.
one amu is exactly 1/12 of the mass of a carbon-12 atom.
Answer
Multiply the fraction of each isotope by its mass and add these numbers to obtain the atomic weight of
magnesium.
Atomic weight = 0.7899(23.98504 amu) + 0.1000(24.98584 amu) + 0.1101(25.98259 amu)
= 18.946 amu + 2.4986 amu + 2.8607 amu
= 24.30 amu (to four significant figures)
The atomic weight of gallium is 69.72 amu. The masses of the naturally occurring isotopes are 68.9257
amu for 6931Ga and 70.9249 amu for 7131Ga. Calculate the percent abundance of each isotope.
Answer
Questions
✓ The atomic weight of lithium is 6.941 amu. The two naturally occurring isotopes of lithium have the
following masses: 6Li, 6.01512 amu; 7Li, 7.01600 amu. Calculate the percent of 6Li in naturally
occurring lithium.
✓ The atomic weight of rubidium is 85.4678 amu. The two naturally occurring isotopes of rubidium
have the following masses: 85Rb, 84.9118 amu; 87Rb, 86.9092 amu. Calculate the percent of 85Rb in
naturally occurring rubidium.
✓ Bromine is composed of 3579Br, 78.9183 amu, and 3581Br 80.9163 amu. The percent composition of
a sample is 50.69% Br-79 and 49.31% Br-81. Based on this sample, calculate the atomic weight of
bromine.
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Course Synopsis:
1. Modern electronic theory of atoms
2. Radioactivity
3. Chemical bonding
MODERN ELECTRONIC THEORY OF ATOMS
In our study of atomic structure, we look first at the fundamental particles. These are the basic building
blocks of all atoms. Atoms, and hence all matter, consist principally of three fundamental particles:
electrons, protons, and neutrons. Knowledge of the nature and functions of these particles is essential to
understanding chemical interactions. The relative masses and charges of the three fundamental particles are
shown in Table
Fundamental Particles of Matter
The mass of an electron is very small compared with the mass of either a proton or a neutron. The charge
on a proton is equal in magnitude, but opposite in sign, to the charge on an electron. The proton is a
fundamental particle with a charge equal in magnitude but opposite in sign to the charge on the electron.
Its mass is almost 1836 times that of the electron.
Note: Many other subatomic particles, such as quarks, positrons, neutrinos, pions, and muons,
have also been discovered.
Rutherford was able to determine the magnitudes of the positive charges on the atomic nuclei. The picture
of atomic structure that he developed is called the Rutherford model of the atom.
Atoms consist of very small, very dense positively charged nuclei surrounded by clouds of
electrons at relatively great distances from the nuclei.
ATOMIC NUMBER
Moseley (1887–1915) showed that the X-ray wavelengths could be better correlated with the atomic
number. On the basis of his mathematical analysis of these X-ray data, he concluded that
each element differs from the preceding element by having one more positive charge in its nucleus.
We now know that every nucleus contains an integral number of protons exactly equal to the number of
electrons in a neutral atom of the element. Every hydrogen atom contains one proton, every helium atom
1|Page
,contains two protons, and every lithium atom contains three protons. The number of protons in the nucleus
of an atom determines its identity; this number is known as the atomic number of that element.
MASS NUMBER AND ISOTOPES
Most elements consist of atoms of different masses, called isotopes. The isotopes of a given element contain
the same number of protons (and also the same number of electrons) because they are atoms of the same
element. They differ in mass because they contain different numbers of neutrons in their nuclei.
Isotopes are atoms of the same element with different masses; they are atoms containing the same
number of protons but different numbers of neutrons.
The mass number of an atom is the sum of the number of protons and the number of neutrons in its
nucleus
Mass number = number of protons + number of neutrons
= atomic number + neutron number
The composition of a nucleus is indicated by its nuclide symbol. This consists of the symbol for the element
(E), with the atomic number (Z) written as a subscript at the lower left and the mass number (A) as a
superscript at the upper left, AZE.
Note of Emphasis
✓ The atomic number, Z, is an integer equal to the number of protons in the nucleus of an atom of the
element. It is also equal to the number of electrons in a neutral atom. It is the same for all atoms of
an element.
✓ The mass number, A, is an integer equal to the sum of the number of protons and the number of
neutrons in the nucleus of an atom of a particular isotope of an element. It is different for different
isotopes of the same element.
✓ Many elements occur in nature as mixtures of isotopes. The atomic weight of such an element is the
weighted average of the masses of its isotopes. Atomic weights are fractional numbers, not integers.
Example:
1. Determine the number of protons, neutrons, and electrons in each of the following species. Are the
members within each pair isotopes?
2|Page
, Calculation of Atomic Weight
Three isotopes of magnesium occur in nature. Their abundances and masses, determined by mass
spectrometry, are listed in the following table below. Use this information to calculate the atomic weight
of magnesium.
one amu is exactly 1/12 of the mass of a carbon-12 atom.
Answer
Multiply the fraction of each isotope by its mass and add these numbers to obtain the atomic weight of
magnesium.
Atomic weight = 0.7899(23.98504 amu) + 0.1000(24.98584 amu) + 0.1101(25.98259 amu)
= 18.946 amu + 2.4986 amu + 2.8607 amu
= 24.30 amu (to four significant figures)
The atomic weight of gallium is 69.72 amu. The masses of the naturally occurring isotopes are 68.9257
amu for 6931Ga and 70.9249 amu for 7131Ga. Calculate the percent abundance of each isotope.
Answer
Questions
✓ The atomic weight of lithium is 6.941 amu. The two naturally occurring isotopes of lithium have the
following masses: 6Li, 6.01512 amu; 7Li, 7.01600 amu. Calculate the percent of 6Li in naturally
occurring lithium.
✓ The atomic weight of rubidium is 85.4678 amu. The two naturally occurring isotopes of rubidium
have the following masses: 85Rb, 84.9118 amu; 87Rb, 86.9092 amu. Calculate the percent of 85Rb in
naturally occurring rubidium.
✓ Bromine is composed of 3579Br, 78.9183 amu, and 3581Br 80.9163 amu. The percent composition of
a sample is 50.69% Br-79 and 49.31% Br-81. Based on this sample, calculate the atomic weight of
bromine.
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