Isotopes and Decay Study Guide

1.2 Isotopes and Decay (Chem.1.2)

Explore this Phenomenon

Some atoms on the chart above are stable, while others are not.

1. What subatomic particle(s) do you predict are the most important in stability?
2. Do unstable atoms follow predictable patterns?

 

Standard Chem.1.2

Analyze and interpret data to identify patterns in the stability of isotopes and predict likely modes of radioactive decay. Emphasize that different isotopes of the same element decay by different modes and at different rates depending on their nuclear stability. Examples of data could include band of stability charts, mass or nuclear binding energy per nucleon, or the inverse relationship between half-life and nuclear stability. (PS1.C)

As you read this section, look for patterns of stability of isotopes. What pattern is there among unstable isotopes? Predict how the nucleus will become stable.

 

History of Atomic Mass Determinations

As a part of his research on atoms, John Dalton determined a number of atomic weights of elements in the early 1800s. Atomic weights were the basis for the periodic table that Mendeleev developed. Originally all atomic weights were based on a comparison to hydrogen, which has an atomic weight of one. After the discovery of the proton, scientists assumed that the weight of an atom was essentially that of the protons – electrons were known to contribute almost nothing to the atomic weight of the element.

This approach worked until we learned how to determine the number of protons in an element . We then saw that the atomic weight for an element was often twice the number of protons (or more). The discovery of the neutron provided the missing part of the picture. The atomic mass is now known to be the sum of the protons and neutrons in the nucleus .

 

Mass Number

Rutherford showed that the vast majority of the mass of an atom is concentrated in its nucleus , which is composed of protons and neutrons. The mass number is defined as the total number of protons and neutrons in an atom . It can be calculated by adding the number of neutrons and the number of protons (atomic number) together.

 

Mass number = atomic number + number of neutrons

 

This table shows the first six elements of the periodic table.

Name	Symbol	Atomic Number	Protons	Neutrons	Electrons	Mass Number
Hydrogen	H	1	1	0	1	1
Helium	He	2	2	2	2	4
Lithium	Li	3	3	4	3	7
Beryllium	Be	4	4	5	4	9
Boron	B	5	5	6	5	11
Carbon	C	6	6	6	6	12

 

Consider the element helium. Its atomic number is 2, so it has two protons in its nucleus . Its nucleus also contains two neutrons. Since 2 + 2 = 4, we know that the mass number of the helium atom is 4. Finally, the helium atom also contains two electrons since the number of electrons must equal the number of protons. This example may lead you to believe that atoms have the same number of protons and neutrons, but further examination of the Table above will show that this is not the case. Lithium, for example has three protons and four neutrons, leaving it with a mass number of 7.

Knowing the mass number and the atomic number of an atom allows you to determine the number of neutrons present in the atom by subtraction.

Number of neutrons = mass number - atomic number

Atoms of the element chromium (Cr) have an atomic number of 24 and a mass number of 52. How many neutrons are in the nucleus of a chromium atom? To determine this, you would subtract as shown:

52 - 24 = 28 neutrons in a chromium atom

The composition of any atom can be illustrated with a shorthand notation using the atomic number and the mass number. Both are written before the chemical symbol, with the mass number written as a superscript and the atomic number written as a subscript. The chromium atom discussed above would be written as:

Another way to refer to a specific atom is to write the mass number of the atom after the name, separated by a hyphen. The above atom would be written as chromium-52.

 

Are all the members of the football team shown above identical?

 

They are on the same team and are all known by the same team name, but there are individual differences among the players. We do not expect the kicker to be as big as the quarterback. The tight end is very likely to weigh less than the defensive tackle on the other side of the ball. They play as a unit, but they have different weights and heights.

 

Isotopes

The history of the atom is full of some of these differences. Although John Dalton stated in his atomic theory of 1804 that all atoms of an element are identical, the discovery of the neutron began to show that this assumption was not correct. The study of radioactive materials (elements that spontaneously give off particles to form new elements) by Frederick Soddy (1877-1956) gave important clues about the internal structure of atoms. His work showed that some substances with different radioactive properties and different atomic masses were in fact the same element . He coined the term isotope from the Greek roots isos (íσος “equal”) and topos (τóπος “place”). He described isotopes as, “Put colloquially, their atoms have identical outsides but different insides.” Soddy won the Nobel Prize in Chemistry in 1921 for his work.

As stated earlier, not all atoms of a given element are identical. Specifically, the number of neutrons can be variable for many elements. As an example, naturally occurring carbon exists in three forms. Each carbon atom has the same number of protons (6), which is its atomic number. Each carbon atom also contains six electrons in order to maintain electrical neutrality. However the number of neutrons varies as six, seven, or eight. Isotopes are atoms that have the same atomic number, but different mass numbers due to a change in the number of neutrons.

The three isotopes of carbon can be referred to as carbon-12 , carbon-13, and carbon-14 refers to the nucleus of a given isotope of an element. A carbon atom is one of three different nuclides. Most elements naturally consist of mixtures of isotopes. Carbon has three natural isotopes, while some heavier elements can have many more. Tin has ten stable isotopes, the most of any element.

While the presence of isotopes affects the mass of an atom, it does not affect its chemical reactivity. Chemical behavior is governed by the number of electrons and the number of protons. Carbon-13 behaves chemically in exactly the same way as the more plentiful carbon-12.

 

Radioactivity

What does this sign mean?

If you visit the nuclear medicine department of a large hospital, you are very likely to see the symbol shown above. The sign means that radioactive materials are present and special safety precautions need to be taken. These materials are used for diagnosis and treatment of many diseases. The people using these materials are specially trained to handle them safely. Radioactive materials can be dangerous and should be respected, but need not be feared.

 

Discovery of Radioactivity

John Dalton first proposed his atomic theory in an 1804 lecture to the Royal Institution, a prestigious British scientific society. In this talk, he put forth the idea that all atoms of an element were identical and that atoms were indestructible. In a little over 100 years, both of these ideas were shown to be incorrect.

In 1919, studies on atomic weights led Francis Aston (1877-1925) to the conclusion that some elements with different atomic weights were actually the same element in different isotope forms. Aston used a mass spectrograph to separate isotopes of different elements. He won the Nobel Prize in Chemistry for this work in 1922.

 

Natural Radioactivity

Pierre and Marie Curie studied the properties of uranium salts with the express purpose of identifying the details of these emissions. They were the first to coin the term “radioactivity,” meaning the spontaneous emission of radiation in the form of particles or high energy photons resulting from a nuclear reaction. The major contributions to the work came from Marie who showed that the amount of radioactivity present was due to the amount of a specific element and not due to some chemical reaction . She discovered the element polonium and named it after her native Poland. Madame Curie shared the 1903 Nobel Prize in Physics with her husband Pierre and Henri Becquerel. She won the Nobel Prize in Chemistry in 1911.

Pierre and Marie Curie in their lab.

 

Ernest Rutherford, a later researcher, was able to show there are three different types of radioactive emissions. These emission types differed in terms of mass, charge, and their ability to penetrate materials. He designated them simply as alpha (α) emissions, beta (β) emissions, and gamma (γ) emissions.

Radioactivity involves the spontaneous emission of material and/or energy from the nucleus of an atom . The most common radioactive atoms have high atomic numbers and contain a large excess of neutrons.

 

Trends in Type of Radioactive Decay

Alpha (α) decay is the emission of an α particle from the nucleus. For example, polonium-210 undergoes α decay:

Alpha decay occurs primarily in heavy nuclei (A > 200, Z > 83). Because the loss of an α particle gives a daughter nuclide with a mass number four units smaller and an atomic number two units smaller than those of the parent nuclide, the daughter nuclide has a larger n:p ratio than the parent nuclide. If the parent nuclide undergoing α decay lies below the band of stability (refer to [link]), the daughter nuclide will lie closer to the band.

Beta (β) decay is the emission of an electron from the nucleus. Iodine-131 is an example of a nuclide that undergoes β decay:

Beta decay, which can be thought of as the conversion of a neutron into a proton and a β particle, is observed in nuclides with a large n:p ratio. The beta particle (electron) emitted is from the atomic nucleus and is not one of the electrons surrounding the nucleus. Such nuclei lie above the band of stability. Emission of an electron does not change the mass number of the nuclide but does increase the number of its protons and decrease the number of its neutrons. Consequently, the n:p ratio is decreased, and the daughter nuclide lies closer to the band of stability than did the parent nuclide.

Gamma emission (γ emission) is observed when a nuclide is formed in an excited state and then decays to its ground state with the emission of a γ ray, a quantum of high-energy electromagnetic radiation. The presence of a nucleus in an excited state is often indicated by an asterisk (*). Cobalt-60 emits γ radiation and is used in many applications including cancer treatment:

There is no change in mass number or atomic number during the emission of a γ ray unless the γ emission accompanies one of the other modes of decay.

Positron emission (β+ decay) is the emission of a positron from the nucleus. Oxygen-15 is an example of a nuclide that undergoes positron emission:

Positron emission is observed for nuclides in which the n:p ratio is low. These nuclides lie below the band of stability. Positron decay is the conversion of a proton into a neutron with the emission of a positron. The n:p ratio increases, and the daughter nuclide lies closer to the band of stability than did the parent nuclide.

 

Putting It Together

Let us revisit the phenomenon:

1. Does the band of nonradioactive elements match up with the straight line that represents a 1:1 ratio of protons to neutrons? Why would that happen?
2. How many protons are in the last nonradioactive element? Are there elements that have a larger number of protons than that?
3. Which would be the radioactive and nonradioactive isotopes of 64 Zn and 60 Zn? Explain how you know.