Isotopes

• What Is The Difference Between Isotopes Of The Same Element
• Isotopes Are Atoms Of The Same Element That Have Different Numbers Of

Isotopes are atoms of the same element that contain different numbers of neutrons. For these species, the number of electrons and protons remain constant. This difference in neutron amount affects the atomic mass (A) but not the atomic number (Z). In a chemical laboratory, isotopes of an element appear and react the same. For this reason, it is difficult to distinguish between an atom's isotopes. In contrast, nuclear scientists can identify and separate different types of atomic nuclei. The technology required for this process is more sophisticated that what could be found in a typical chemical laboratory.

Atoms of the same element with different numbers of neutrons are called isotopes. Isotopes of an element have: the same atomic number. Different mass numbers. Three isotopes of hydrogen. Hydrogen (1 H) has three naturally occurring isotopes, sometimes denoted 1 H, 2 H, and 3 H. The first two of these are stable, while 3 H has a half-life of 12.32 years. There are also heavier isotopes, which are all synthetic and have a half-life less than one zeptosecond (10 −21 second). Of these, 5 H is the most stable, and 7 H is the least. Hydrogen is the only element whose isotopes have. What are isotopes? A.) Atoms of the same element that have a different number of protons B.) Atoms of the same element that have a different number of neutrons C.) Atoms of different elements that have the same number of neutrons D.) Atoms that are unstable.

The element carbon ((ce{C})) has an atomic number of 6, which means that all neutral carbon atoms contain 6 protons and 6 electrons. In a typical sample of carbon-containing material, 98.89% of the carbon atoms also contain 6 neutrons, so each has a mass number of 12. An isotope of any element can be uniquely represented as ({}_Z^{A}X) where X is the atomic symbol of the element. The isotope of carbon that has 6 neutrons is therefore (ce{_6^{12}C}) The subscript indicating the atomic number is actually redundant because the atomic symbol already uniquely specifies Z. Consequently, ​it is more often written as (ce{^{12}C}), which is read as “carbon-12.” Nevertheless, the value of (Z) is commonly included in the notation for nuclear reactions because these reactions involve changes in (Z).

Most elements on the periodic table have at least two stable isotopes. For example, in addition to (ce{^{12}C}), a typical sample of carbon contains 1.11% (ce{_6^{13}C}), with 7 neutrons and 6 protons, and a trace of (ce{_6^{14}C}), with 8 neutrons and 6 protons. The nucleus of (ce{_6^{14}C}) is not stable, however, but undergoes a slow radioactive decay that is the basis of the carbon-14 dating technique used in archeology. Many elements other than carbon have more than one stable isotope; tin, for example, has 10 isotopes. There are about twenty elements that exist in only one isotopic form (sodium and fluorine are examples of these).

An important series of isotopes is found with hydrogen atoms. Most hydrogen atoms have a nucleus with only a single proton. About 1 in 10,000 hydrogen nuclei, however, also has a neutron; this particular isotope is called deuterium. An extremely rare hydrogen isotope, tritium, has 1 proton and 2 neutrons in its nucleus. Figure (PageIndex{1}) compares the three isotopes of hydrogen.

There are currently over 3,500 isotopes known for all the elements. When scientists discuss individual isotopes, they need an efficient way to specify the number of neutrons in any particular nucleus. A/Z and symbol-mass formats can be used to display periodic table information. When viewing either of these two notations, isotopic differences can be obtained.

The discovery of isotopes required a minor change in Dalton’s atomic theory. Dalton thought that all atoms of the same element were exactly the same.

Look at the A/Z formats for the three isotopes of hydrogen in Table (PageIndex{1}). Note how the atomic number (bottom value) remains the same while the atomic masses (top number) are varied. All isotopes of a particular element will vary in neutrons and mass. This variance in mass will be visible in the symbol-mass format of same isotopes as well.

Both A/Z or symbol-mass formats can be utilized to determine the amount of subatomic particles (protons, neutrons, and electrons) contained inside an isotope. When given either format, these mass values should be used to calculate the number of neutrons in the nucleus.

GENERAL CHEMISTRY TOPICS

Isotopes

An isotope is made up of atoms of the same element that have the same atomic mass. Different isotopes of an element arise from atoms with differing numbers of neutrons. Uses of isotopes. Average atomic masses from natural abundances: The weighted-average calculation.

Atomic number, mass number and isotopes

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The atomic number of an element (symbolized as Z) is the number of protons in the nuclei of its atoms. The mass number (A) is the total number of nucleons (neutrons and protons). An isotope is made up of atoms of the same element (which by definition have a characteristic and fixed atomic number) that also have the same mass number. Different isotopes of an element arise from atoms with differing numbers of neutrons. Because of this, chemists need a way to represent specific isotopes of an element. Isotopes of an element have the same atomic number, but different mass numbers. The atomic number, when represented along with the symbol of an element, is shown as a leading subscript. The mass number is shown as a leading superscript. Since the element symbol implies an atomic number, the latter is often dropped, and an isotope as commonly represented textually with just the mass number and the element symbol (for example ¹⁴C or ¹⁸O).

In the periodic table, the elements, represented as their symbols, are arranged in a particular pattern that reflects (as we will see) a regularity, or periodicity in their properties. Typically in the table, the element symbol is contained within its own small box, along with other information including the atomic number and the average atomic mass. The average atomic mass of an element represents the averages of its naturally occurring isotopic masses weighted according to their natural abundance. The formula for calculation of average atomic mass and illustration of its use is presented below.

What Is The Difference Between Isotopes Of The Same Element

How do the isotopic forms of an element differ from one another, physically and chemically? Isotopes are defined by their subatomic particle composition, which we will think of as a physical property. The chemistry of an element is determined by, in a general sense, the number of valence electrons its atoms possess. The number of valence electrons associated with a neutral atom is in turn determined by the number protons in the nucleus. Thus, two atomic nuclei could have the same number of protons, but different numbers of neutrons. Yet since the atoms they are part of would still have the same number of valence electrons, these two atoms would be chemically indistinguishable.*

Uses of isotopes

There are a wide variety of applications of isotopes in nuclear chemistry, medicine, biochemistry, anthropology, paleontology, and geology. Many such uses are based on the phenomenon of radioactivity, shown by some of the isotopes of many of the elements. Such radioactive isotopes are unstable, undergoing spontaneous nuclear decay processes at a rate determined by the half-life of the isotope. One example is the use of ¹⁴C - the isotope of carbon with six protons and eight neutrons, which has a half-life of 5730 years - as a basis for dating of materials derived from living organisms that are many thousands of years old. This technique, called radiocarbon dating, is used widely in geosciences and anthropology.

Average atomic masses from natural abundances: The weighted-average calculation

The atomic masses given in the periodic table represent weighted averages based on the natural abundances of the isotopes of a given element. The formula for a weighted average is

Here the x_i's are the masses of the individual isotopes, and the w_i's are the fractional abundances corresponding to the isotopes. Note that the weights must sum to 1 (equivalently the percent abundances must sum to 100%).

For example, chlorine exists in two isotopic forms, ³⁵Cl and ³⁷Cl. The mass of the ³⁵Cl isotope is 34.97 amu and that of ³⁷Cl is 36.97 amu. The abundances are 75.77% and 24.23%, repectively. Therefore in this case, the weighted average becomes

wa = (0.7577)(34.97 amu) + (0.2423)(36.97 amu) = 35.45 amu

The result of this calculation is the atomic mass of chlorine that appears in the periodic table.

Isotopes Are Atoms Of The Same Element That Have Different Numbers Of

* Actually, since isotopes of an element differ in atomic mass, they can be subtly distinguished by differences in reaction rates, or in physical processes - such as rate of diffusion - affected by mass.