# 1.8: Subatomic Particles - Protons, Neutrons, and Electrons


Learning Objectives

• To know the meaning of isotopes and atomic masses.

To date, about 118 different elements have been discovered; by definition, each is chemically unique. To understand why they are unique, you need to understand the structure of the atom (the fundamental, individual particle of an element) and the characteristics of its components. Atoms consist of electrons, protons, and neutrons. Although this is an oversimplification that ignores the other subatomic particles that have been discovered, it is sufficient for discussion of chemical principles. Some properties of these subatomic particles are summarized in Table $$\PageIndex{1}$$, which illustrates three important points:

1. Electrons and protons have electrical charges that are identical in magnitude but opposite in sign. Relative charges of −1 and +1 are assigned to the electron and proton, respectively.
2. Neutrons have approximately the same mass as protons but no charge. They are electrically neutral.
3. The mass of a proton or a neutron is about 1836 times greater than the mass of an electron. Protons and neutrons constitute the bulk of the mass of atoms.

The discovery of the electron and the proton was crucial to the development of the modern model of the atom and provides an excellent case study in the application of the scientific method. In fact, the elucidation of the atom’s structure is one of the greatest detective stories in the history of science.

Table $$\PageIndex{1}$$: Properties of Subatomic Particles*
Particle Mass (g) Atomic Mass (amu) Electrical Charge (coulombs) Relative Charge
Location
electron $$9.109 \times 10^{-28}$$ 0.0005486 −1.602 × 10−19 −1
Outside nucleus
proton $$1.673 \times 10^{-24}$$ 1.007276 +1.602 × 10−19 +1
In nucleus
neutron $$1.675 \times 10^{-24}$$ 1.008665 0 0
In nucleus

Almost all of the mass of an atom is contained within a tiny (and therefore extremely dense) nucleus which carries a positive electric charge and almost all of the volume of an atom consists of empty space in which electrons reside (Figure $$\PageIndex{1}$$). The extremely small mass of the electron (1/1840 the mass of the hydrogen nucleus) causes it to behave as a quantum particle, which means that its location at any moment cannot be specified; the best we can do is describe its behavior in terms of the probability of its manifesting itself at any point in space. It is common (but somewhat misleading) to describe the volume of space in which the electrons of an atom have a significant probability of being found as the electron cloud. The latter has no definite outer boundary, so neither does the atom. The radius of an atom must be defined arbitrarily, such as the boundary in which the electron can be found with 95% probability. Atomic radii are typically 30-300 pm.

## The Number of Protons Define the Nature of the Elements

The number of protons in the nucleus of an atom is its atomic number ($$Z$$). This is the defining trait of an element: Its value determines the identity of the atom. For example, any atom that contains six protons is the element carbon and has the atomic number 6, regardless of how many neutrons or electrons it may have. A neutral atom must contain the same number of positive and negative charges, so the number of protons equals the number of electrons. Therefore, the atomic number also indicates the number of electrons in an atom. The total number of protons and neutrons in an atom is called its mass number ($$A$$)). The number of neutrons is therefore the difference between the mass number and the atomic number:

\begin{align*} \ce{atomic\: number\:(Z)\: &= \:number\: of\: protons\\ mass\: number\:(A)\: &= \:number\: of\: protons + number\: of\: neutrons\\ A-Z\: &= \:number\: of\: neutrons} \end{align*}

The identity of an element is defined by $$Z$$, the number of protons in the nucleus of an atom of the element, which is different for each element. The known elements are arranged in order of increasing Z in the periodic table (Figure $$\PageIndex{2}$$). The rationale for the peculiar format of the periodic table is explained later. Each element is assigned a unique one-, two-, or three-letter symbol. The names of the elements are listed in the periodic table, along with their symbols, atomic numbers, and atomic masses. The chemistry of each element is determined by its number of protons and electrons.

The symbols for several common elements and their atoms are listed in Table $$\PageIndex{2}$$. Some symbols are derived from the common name of the element; others are abbreviations of the name in another language. Symbols have one or two letters, for example, H for hydrogen and $$\ce{Cl}$$ for chlorine. To avoid confusion with other notations, only the first letter of a symbol is capitalized. For example, $$\ce{Co}$$ is the symbol for the element cobalt, but $$\ce{CO}$$ is the notation for the compound carbon monoxide, which contains atoms of the elements carbon ($$\ce{C}$$) and oxygen ($$\ce{O}$$). All known elements and their symbols are in the periodic table.

Table $$\PageIndex{2}$$: Some Common Elements and Their Symbols
Element Symbol Element Symbol
aluminum Al iron Fe (from ferrum)
bromine Br lead Pb (from plumbum)
calcium Ca magnesium Mg
carbon C mercury Hg (from hydrargyrum)
chlorine Cl nitrogen N
chromium Cr oxygen O
cobalt Co potassium K (from kalium)
copper Cu (from cuprum) silicon Si
fluorine F silver Ag (from argentum)
gold Au (from aurum) sodium Na (from natrium)
helium He sulfur S
hydrogen H tin Sn (from stannum)
iodine I zinc Zn

Traditionally, the discoverer (or discoverers) of a new element names the element. However, until the name is recognized by the International Union of Pure and Applied Chemistry (IUPAC), the recommended name of the new element is based on the Latin word(s) for its atomic number. For example, element 106 was called unnilhexium (Unh), element 107 was called unnilseptium (Uns), and element 108 was called unniloctium (Uno) for several years. These elements are now named after scientists or locations; for example, element 106 is now known as seaborgium (Sg) in honor of Glenn Seaborg, a Nobel Prize winner who was active in the discovery of several heavy elements

## Isotopes: Differing Numbers of Neutrons

Recall that the nuclei of most atoms contain neutrons as well as protons. Unlike protons, the number of neutrons is not absolutely fixed for most elements. Atoms that have the same number of protons, and hence the same atomic number, but different numbers of neutrons are called isotopes. All isotopes of an element have the same number of protons and electrons, which means they exhibit the same chemistry. The isotopes of an element differ only in their atomic mass, which is given by the mass number ($$A$$), the sum of the numbers of protons and neutrons.

The element carbon ($$C$$) 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 $$^A_Z X$$, where X is the atomic symbol of the element. The isotope of carbon that has 6 neutrons is therefore $$_6^{12} C$$. The subscript indicating the atomic number is actually redundant because the atomic symbol already uniquely specifies Z. Consequently, $$_6^{12} C$$ is more often written as 12C, 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.

For example, naturally occurring hydrogen has two stable nuclides, $$\ce{^{1}_{1}H}$$ and $$\ce{^{2}_{1}H}$$, which also are isotopes of one another. More than 99.98 percent is “light” hydrogen, $$\ce{^{1}_{1}H}$$. This consists of atoms each of which has one proton, one electron, and zero neutrons (Figure $$\PageIndex{1; left}$$). The rest is “heavy” hydrogen or deuterium, $$\ce{^{2}_{1}H}$$, which consists of atoms which contain one electron, one proton, and one neutron (Figure $$\PageIndex{1; center}$$. Hence the nuclidic mass of deuterium is almost exactly twice as great as for light hydrogen. It is also possible to obtain a third isotope, tritium, $$\ce{^{3}_{1}H}$$. that consists of atoms whose nuclei contain two neutrons and one proton (Figure $$\PageIndex{1; right}$$). Its mass is about 3 times that of light hydrogen.

In addition to $$^{12}C$$, a typical sample of carbon contains 1.11% $$_6^{13} C$$ (13C), with 7 neutrons and 6 protons, and a trace of $$_6^{14} C$$ (14C), with 8 neutrons and 6 protons. The nucleus of 14C 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. Information about the naturally occurring isotopes of elements with atomic numbers 1 through 10 is given in Table $$\PageIndex{2}$$. Note that in addition to standard names and symbols, the isotopes of hydrogen are often referred to using common names and accompanying symbols. Hydrogen-2, symbolized 2H, is also called deuterium and sometimes symbolized D. Hydrogen-3, symbolized 3H, is also called tritium and sometimes symbolized T.

Table $$\PageIndex{3}$$: Nuclear Compositions of Atoms of the Very Light Elements
Element Symbol Atomic Number Number of Protons Number of Neutrons Mass (amu) % Natural Abundance
hydrogen $$\ce{^1_1H}$$
(protium)
1 1 0 1.0078 99.989
$$\ce{^2_1H}$$
(deuterium)
1 1 1 2.0141 0.0115
$$\ce{^3_1H}$$
(tritium)
1 1 2 3.01605 — (trace)
helium $$\ce{^3_2He}$$ 2 2 1 3.01603 0.00013
$$\ce{^4_2He}$$ 2 2 2 4.0026 100
lithium $$\ce{^6_3Li}$$ 3 3 3 6.0151 7.59
$$\ce{^7_3Li}$$ 3 3 4 7.0160 92.41
beryllium $$\ce{^9_4Be}$$ 4 4 5 9.0122 100
boron $$\ce{^{10}_5B}$$ 5 5 5 10.0129 19.9
$$\ce{^{11}_5B}$$ 5 5 6 11.0093 80.1
carbon $$\ce{^{12}_6C}$$ 6 6 6 12.0000 98.89
$$\ce{^{13}_6C}$$ 6 6 7 13.0034 1.11
$$\ce{^{14}_6C}$$ 6 6 8 14.0032 — (trace)
nitrogen $$\ce{^{14}_7N}$$ 7 7 7 14.0031 99.63
$$\ce{^{15}_7N}$$ 7 7 8 15.0001 0.37
oxygen $$\ce{^{16}_8O}$$ 8 8 8 15.9949 99.757
$$\ce{^{17}_8O}$$ 8 8 9 16.9991 0.038
$$\ce{^{18}_8O}$$ 8 8 10 17.9992 0.205
fluorine $$\ce{^{19}_9F}$$ 9 9 10 18.9984 100
neon $$\ce{^{20}_{10}Ne}$$ 10 10 10 19.9924 90.48
$$\ce{^{21}_{10}Ne}$$ 10 10 11 20.9938 0.27
$$\ce{^{22}_{10}Ne}$$ 10 10 12 21.9914 9.2

Sources of isotope data: G. Audi et al., Nuclear Physics A 729 (2003): 337–676; J. C. Kotz and K. F. Purcell, Chemistry and Chemical Reactivity, 2nd ed., 1991.

Example $$\PageIndex{1}$$: Composition of an Atom

Iodine is an essential trace element in our diet; it is needed to produce thyroid hormone. Insufficient iodine in the diet can lead to the development of a goiter, an enlargement of the thyroid gland. The addition of small amounts of iodine to table salt (iodized salt) has essentially eliminated this health concern in the United States, but as much as 40% of the world’s population is still at risk of iodine deficiency. The iodine atoms are added as anions, and each has a 1− charge and a mass number of 127. Determine the numbers of protons, neutrons, and electrons in one of these iodine anions.

Solution

The atomic number of iodine (53) tells us that a neutral iodine atom contains 53 protons in its nucleus and 53 electrons outside its nucleus. Because the sum of the numbers of protons and neutrons equals the mass number, 127, the number of neutrons is 74 (127 − 53 = 74). Since the iodine is added as a 1− anion, the number of electrons is 54 [53 – (1–) = 54].

Exercise $$\PageIndex{1}$$

An ion of platinum has a mass number of 195 and contains 74 electrons. How many protons and neutrons does it contain, and what is its charge?

78 protons; 117 neutrons; charge is 4+

Example $$\PageIndex{2}$$

An element with three stable isotopes has 82 protons. The separate isotopes contain 124, 125, and 126 neutrons. Identify the element and write symbols for the isotopes.

Given: number of protons and neutrons

Asked for: element and atomic symbol

Strategy:

1. Refer to the periodic table and use the number of protons to identify the element.
2. Calculate the mass number of each isotope by adding together the numbers of protons and neutrons.
3. Give the symbol of each isotope with the mass number as the superscript and the number of protons as the subscript, both written to the left of the symbol of the element.

Solution:

A The element with 82 protons (atomic number of 82) is lead: Pb.

B For the first isotope, A = 82 protons + 124 neutrons = 206. Similarly, A = 82 + 125 = 207 and A = 82 + 126 = 208 for the second and third isotopes, respectively. The symbols for these isotopes are $$^{206}_{82}Pb$$, $$^{207}_{82}Pb$$, and $$^{208}_{82}Pb$$, which are usually abbreviated as $$^{206}Pb$$, $$^{207}Pb$$, and $$^{208}Pb$$.

Exercise $$\PageIndex{2}$$

Identify the element with 35 protons and write the symbols for its isotopes with 44 and 46 neutrons.

$$\ce{^{79}_{35}Br}$$ and $$\ce{^{81}_{35}Br}$$ or, more commonly, $$\ce{^{79}Br}$$ and $$\ce{^{81}Br}$$.

## Ions: Charged

The protons and neutrons in the nucleus of an atom are held very tightly by strong nuclear forces. It is very difficult either to separate the nuclear particles or to add extra ones. The electrons, on the other hand, are held to the atom by their electrostatic attraction for the positively charged protons in the nucleus. This force is strong, but not so strong that an atom cannot lose or gain electrons. Atoms are electrically neutral if they contain the same number of positively charged protons and negatively charged electrons. When the numbers of these subatomic particles are not equal, the atom is electrically charged and is called an ion. The charge of an atom is defined as follows:

$\text{Atomic charge} = \text{number of protons} − \text{number of electrons}$

As will be discussed in more detail later in this chapter, atoms (and molecules) typically acquire charge by gaining or losing electrons. An atom that gains one or more electrons will exhibit a negative charge and is called an anion. Positively charged atoms called cations are formed when an atom loses one or more electrons. For example, a neutral sodium atom (Z = 11) has 11 electrons. If this atom loses one electron, it will become a cation with a 1+ charge (11 − 10 = 1+).

$\ce{Na + e^{-} -> Na^{+}}$

Similarly, a neutral fluorine atom (F = 9) has nine electrons, and if it gains a electron it will become an anion with a 1− charge (9 − 10 = 1−).

$\ce{F - 2e^{-} -> F^{-2}}$

The charge of the species has a profound affect on the properties of the species. For example, neutral sodium atoms are unstable and reacting violently when combined with most substnces. However, sodium cations are quite inert; in fact, we eat them all the time as part of sodium chloride (table salt). As we will discussed later, cations and anions almost always occur together to ensure that matter is neutral.

## Summary

The atom consists of discrete particles that govern its chemical and physical behavior. Each atom of an element contains the same number of protons, which is the atomic number (Z). Neutral atoms have the same number of electrons and protons. Atoms of an element that contain different numbers of neutrons are called isotopes. Each isotope of a given element has the same atomic number but a different mass number (A), which is the sum of the numbers of protons and neutrons. The relative masses of atoms are reported using the atomic mass unit (amu), which is defined as one-twelfth of the mass of one atom of carbon-12, with 6 protons, 6 neutrons, and 6 electrons.