1.3 Atomic Mass Atomic Number and Isotopes PDF

Atomic Number, Mass
Number, and Isotopes

Prepared by:
Engr. Adrian R. Agunday, rMP, LPT
Special Science Teacher 1
Atomic Number and Mass Number
Atomic number (Z)
- is the number of protons in the nucleus
of each atom of an element.
- In a neutral atom the number of protons
is equal to the number of electrons, but
for those atoms that are not neutral
(with charge) the number of protons and
electrons are not equal.
- The chemical identity of an atom can be
determined solely from its atomic
number.
For example, the atomic number of fluorine is 9. This means that
each fluorine atom has 9 protons and 9 electrons. Or, viewed another way,
every atom in the universe that contains 9 protons is correctly named
“fluorine.”
Mass number (A)

- is the total number of neutrons and protons present in the nucleus of an
atom of an element.
- Except for the most common form of hydrogen, which has one proton and
no neutrons
• The number of neutrons in an atom is equal to the difference between the
mass number and the atomic number, or (A - Z).
For example, if the mass number of a

particular boron atom is 12 and the atomic
number is 5 (indicating 5 protons in the
nucleus), then the number of neutrons is 12 -
5 = 7.
Isotopes
Atoms of a given element do not all have the same mass. Most elements
have two or more isotopes, atoms that have the same atomic number but
different mass numbers.
The Three Hydrogen Isotopes
ISOTOPE PROTIUM DEUTERIUM TRITIUM

Atomic No. 1 1 1
Mass No. 1 2 3
No. of p+ 1 1 1
No. of e- 1 1 1
No. of n 0 1 2
Isotopes
Isotopes of Carbon
Note:
The chemical properties of an element are determined primarily by
the protons and electrons in its atoms; neutrons do not take part in
chemical changes under normal conditions. Therefore, isotopes of the same
element have similar chemistries, forming the same types of compounds and
displaying similar reactivities.
Comparing the Atomic Mass, Atomic Number, No. of protons,
neutrons and electrons of a neutral atom and atom with charge
(positive or negative)
Neutral Atom: Charged Atom:

𝟏𝟎𝟖 𝟏𝟎𝟖 𝟐−
𝟒𝟕𝑨𝒈 𝟒𝟕 𝑨𝒈
Name of Element: Name of Element:
Atomic No.: Atomic No.:
Mass No.: Mass No.:
No. of Protons: No. of Protons:
No. of Neutrons: No. of Neutrons:
No. of Electrons: No. of Electrons:
Charge: Charge:
Note:
The charge of the atom of an element affects the number of protons
and electrons such as:
• If the charge of atom is positive, this means that the number of protons
is greater compared to the number of electrons by the numerical value
of its charge.
Example:
𝟏𝟑𝟕 𝟏+
𝟓𝟔 𝑩𝒂 ,
the number of protons of Barium is 56 and since the
charge is 1+ the number of protons is greater than the number of
electrons by 1, thus the number of electrons is 55.
• If the charge of atom is negative, this means that the number of
protons is lesser compared to the number of electrons by the numerical
value of its charge.
Example:
𝟏𝟑𝟕 𝟏− ,
𝟓𝟔 𝑩𝒂 the number of protons of Barium is 56 and since the
charge is 1- the number of protons is lesser than the number of electrons
by 1, thus the number of electrons is 57.
Example:
Give the number of protons, neutrons, and electrons in each of the
following species:
a. 20
11𝑁𝑎
𝟖𝟎 𝟏−
b. 𝟑𝟓 𝑩𝒓
𝟏𝟏𝟐 𝟐+
c. 𝟒𝟖 𝑪𝒅
𝟏𝟐𝟖 𝟐−
d. 𝟓𝟐 𝑻𝒆
𝟐𝟎𝟏 𝟐+
e. 𝟖𝟎 𝑯𝒈
𝟏𝟖𝟒 𝟓−
f. 𝟕𝟒𝑾
Atomic Mass
The mass of an atom depends on the number
of electrons, protons, and neutrons it contains.
Knowledge of an atom’s mass is important in
laboratory work. But atoms are extremely small
particles—even the smallest speck of dust that our
unaided eyes can detect contains as many as 1 × 1016
atoms!
By international agreement, atomic mass

(sometimes called atomic weight ) is the mass of the
atom in atomic mass units (amu). One atomic mass
unit is defined as a mass exactly equal to one-twelfth
the mass of one carbon-12 atom.
Average Atomic Mass
When you look up the atomic mass of carbon in a table, you will find
that its value is not 12.00 amu but 12.01 amu.
Why?
Average Atomic Mass
The reason for the difference is that most naturally occurring

elements (including carbon) have more than one isotope. This means that
when we measure the atomic mass of an element, we must generally settle for
the average mass of the naturally occurring mixture of isotopes.
For example, the natural abundances of carbon-12 are 98.90 percent
and carbon-13 are 1.10 percent, respectively. The atomic mass of carbon-13
has been determined to be 13.00335 amu. Thus, the average atomic mass of
carbon can be calculated as follows:
Average Atomic Mass of Natural Carbon

= (0.9890)(12.00000 amu) + (0.0110)(13.00335 amu)
= 12.01 amu
It is important to understand that when we say that the atomic mass of
carbon is 12.01 amu, we are referring to the average value. If carbon atoms
could be examined individually, we would find either an atom of atomic mass
12.00000 amu or one of 13.00335 amu, but never one of 12.01 amu.
Try to compute the average
atomic mass!
The atomic masses of its two

stable isotopes, 63
29𝐶𝑢 (69.09 percent)
and 65
29𝐶𝑢 (30.91 percent), are 62.93
amu and 64.9278 amu, respectively.
Calculate the average atomic mass of
copper. The relative abundances are
given in parentheses.
Try to compute the average
atomic mass!
The atomic masses of its two

stable isotopes, 63
29𝐶𝑢 (69.09 percent)
and 65
29𝐶𝑢 (30.91 percent), are 62.93
amu and 64.9278 amu, respectively.
Calculate the average atomic mass of
copper. The relative abundances are
given in parentheses.
Answer: 63.55 amu
Assignment (1 whole sheet of paper)
1. The atomic masses of the two stable isotopes of boron, 105𝐵 (19.78
percent) and 115𝐵 (80.22 percent), are 10.0129 amu and 11.0093 amu,
respectively. Calculate the average atomic mass of boron.
2. The atomic masses of 35

17 𝐶𝑙 (75.53 percent) and 37
17𝐶𝑙 (24.47 percent) are
34.968 amu and 36.956 amu, respectively. Calculate the average atomic
mass of chlorine. The percentages in parentheses denote the relative
abundances.

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Atomic Number, Mass

Number, and Isotopes

Mass number (A)

For example, if the mass number of a

The Three Hydrogen Isotopes

ISOTOPE PROTIUM DEUTERIUM TRITIUM

Neutral Atom: Charged Atom:

By international agreement, atomic mass

The reason for the difference is that most naturally occurring

Average Atomic Mass of Natural Carbon

The atomic masses of its two

The atomic masses of its two

2. The atomic masses of 35

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