The development of modern atomic theory revealed much about the inner structure of atoms. It was learned that an atom contains a very small nucleus composed of positively charged protons and uncharged neutrons, surrounded by a much larger volume of space containing negatively charged electrons. The nucleus contains the majority of an atom’s mass because protons and neutrons are much heavier than electrons, whereas electrons occupy almost all of an atom’s volume. The diameter of an atom is on the order of 10⁻¹⁰ m, whereas the diameter of the nucleus is roughly 10⁻¹⁵ m—about 100,000 times smaller. For a perspective about their relative sizes, consider this: If the nucleus were the size of a blueberry, the atom would be about the size of a football stadium (Figure 1)!

Atoms—and the protons, neutrons, and electrons that compose them—are extremely small. For example, a carbon atom weighs less than 2 × 10⁻²³ g, and an electron has a charge of less than 2 × 10⁻¹⁹ C (coulomb). When describing the properties of tiny objects such as atoms, we use appropriately small units of measure, such as the atomic mass unit (amu) and the fundamental unit of charge (e). The amu was originally defined based on hydrogen, the lightest element, but now is defined with regard to the most abundant isotope of carbon, atoms of which are assigned masses of exactly 12 amu. The fundamental unit of charge (also called the elementary charge) equals the magnitude of the charge of an electron (e) with e = 1.602 × 10⁻¹⁹ C.

A proton has a mass of 1.0073 amu and a charge of 1+. A neutron is a slightly heavier particle with a mass 1.0087 amu and a charge of zero; as its name suggests, it is neutral. The electron has a charge of 1− and is a much lighter particle with a mass of about 0.00055 amu (it would take about 1800 electrons to equal the mass of one proton.

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. 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: A – Z = number of neutrons.

[latex]\begin{array}{r @ {{}={}} l} \text{atomic number (Z)} & \text{number of protons} \\[1em] \text{mass number (A)} & \text{number of protons + number of neutrons} \\[1em] \text{A - Z} & \text{number of neutrons} \end{array}[/latex]

Atoms are electrically neutral with 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.

Atomic or ionic charge = number of protons − number of 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+). A neutral oxygen atom (Z = 8) has eight electrons, and if it gains two electrons it will become an anion with a 2− charge (8 − 10 = 2−).

Chemical Symbols

A chemical symbol is an abbreviation that we use to indicate an element or an atom of an element. For example, the symbol for mercury is Hg (Figure 3). We use the same symbol to indicate one atom of mercury (microscopic domain) or to label a container of many atoms of the element mercury (macroscopic domain).

The symbols for several common elements and their atoms are listed in Table 2. To avoid confusion with other notations, only the first letter of a symbol is capitalized. All known elements and their symbols are in the periodic table in Figure 2 in Chapter 3.5 The Periodic Table (also found in Appendix A).

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
Table 2. Some Common Elements and Their Symbols

Visit this site to learn more about IUPAC, the International Union of Pure and Applied Chemistry, which governs chemical nomenclature (naming). At the IUPAC site you may enjoy exploring their periodic table.

 

Isotopes

The symbol for a specific isotope of any element is written by placing the mass number as a superscript to the left of the element symbol (Figure 4). The atomic number is sometimes written as a subscript preceding the symbol, but since this number defines the element’s identity, as does its symbol, it is often omitted. For example, magnesium exists as a mixture of three isotopes, each with an atomic number of 12 and with mass numbers of 24, 25, and 26, respectively. These isotopes can be identified as ²⁴Mg, ²⁵Mg, and ²⁶Mg. These isotope symbols are read as “element, mass number” and can be symbolized consistent with this reading. For instance, ²⁴Mg is read as “magnesium 24,” and can be written as “magnesium-24” or “Mg-24.” ²⁵Mg is read as “magnesium 25,” and can be written as “magnesium-25” or “Mg-25.” All magnesium atoms have 12 protons in their nucleus. They differ only because a ²⁴Mg atom has 12 neutrons in its nucleus, a ²⁵Mg atom has 13 neutrons, and a ²⁶Mg has 14 neutrons.  Therefore the masses of isotopes of an element, isotopic mass, differ – see Table 3.

 

Information about the naturally occurring isotopes of elements with atomic numbers 1 through 10 is given in Table 3. 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.

Element	Symbol	Atomic Number	# of Protons	# of Neutrons	Isotopic Mass (amu)	% Natural Abundance
hydrogen	[latex]_1^1\text{H}[/latex] (protium)	1	1	0	1.0078	99.989
	[latex]_1^2\text{H}[/latex] (deuterium)	1	1	1	2.0141	0.0115
	[latex]_1^3\text{H}[/latex] (tritium)	1	1	2	3.01605	— (trace)
helium	[latex]_2^3\text{He}[/latex]	2	2	1	3.01603	0.00013
	[latex]_2^4\text{He}[/latex]	2	2	2	4.0026	100
lithium	[latex]_3^6\text{Li}[/latex]	3	3	3	6.0151	7.59
	[latex]_3^7\text{Li}[/latex]	3	3	4	7.0160	92.41
beryllium	[latex]_4^9\text{Be}[/latex]	4	4	5	9.0122	100
boron	[latex]_5^{10}\text{B}[/latex]	5	5	5	10.0129	19.9
	[latex]_5^{11}\text{B}[/latex]	5	5	6	11.0093	80.1
carbon	[latex]_6^{12}\text{C}[/latex]	6	6	6	12.0000	98.89
	[latex]_6^{13}\text{C}[/latex]	6	6	7	13.0034	1.11
	[latex]_6^{14}\text{C}[/latex]	6	6	8	14.0032	— (trace)
nitrogen	[latex]_7^{14}\text{N}[/latex]	7	7	7	14.0031	99.63
	[latex]_7^{15}\text{N}[/latex]	7	7	8	15.0001	0.37
oxygen	[latex]_8^{16}\text{O}[/latex]	8	8	8	15.9949	99.757
	[latex]_8^{17}\text{O}[/latex]	8	8	9	16.9991	0.038
	[latex]_8^{18}\text{O}[/latex]	8	8	10	17.9992	0.205
fluorine	[latex]_9^{19}\text{F}[/latex]	9	9	10	18.9984	100
neon	[latex]_{10}^{20}\text{Ne}[/latex]	10	10	10	19.9924	90.48
	[latex]_{10}^{21}\text{Ne}[/latex]	10	10	11	20.9938	0.27
	[latex]_{10}^{22}\text{Ne}[/latex]	10	10	12	21.9914	9.25
Table 3. Nuclear Compositions of Atoms of the Very Light Elements

Example 3

1. What is the symbol for an isotope of uranium that has an atomic number of 92 and a mass number of 235?
2. How many protons and neutrons are in [latex]_{26}^{56}\text{Fe}[/latex]?

 

Solution

1. The symbol for this isotope is [latex]_{92}^{235}\text{U}[/latex].
2. This iron atom has 26 protons and 56 − 26 = 30 neutrons.

 

Test Yourself

How many protons are in [latex]_{23}^{11}\text{Na}[/latex]?

 

Answer

11 protons

Atomic Mass

Because each proton and each neutron contribute approximately one amu to the mass of an atom, and each electron contributes far less, the atomic mass of a single atom is approximately equal to its mass number (a whole number). However, the average masses of atoms of most elements are not whole numbers because most elements exist naturally as mixtures of two or more isotopes.

The mass of an element shown in a periodic table or listed in a table of atomic masses is a weighted, average mass of all the isotopes present in a naturally occurring sample of that element. This is equal to the sum of each individual isotope’s mass multiplied by its fractional abundance.

[latex]\displaystyle{} \text{average mass} = \sum_{i} (\text{fractional abundance} \times \text{isotopic mass})_{i}[/latex]

Example 6

A meteorite found in central Indiana contains traces of the noble gas neon picked up from the solar wind during the meteorite’s trip through the solar system. Analysis of a sample of the gas showed that it consisted of 91.84% ²⁰Ne (mass 19.9924 amu), 0.47% ²¹Ne (mass 20.9940 amu), and 7.69% ²²Ne (mass 21.9914 amu). What is the average mass of the neon in the solar wind?

 

Solution

[latex]\begin{array}{r @{{}={}} l} \text{average mass} & (0.9184 \times 19.9924 \;\text{amu}) + (0.0047 \times 20.9940 \;\text{amu})+(0.0769 \times 21.9914 \;\text{amu}) \\[1em] & (18.36+0.099+1.69) \;\text{amu} \\[1em] & 20.15 \;\text{amu} \end{array}[/latex]

The average mass of a neon atom in the solar wind is 20.15 amu. (The average mass of a terrestrial neon atom is 20.1796 amu. This result demonstrates that we may find slight differences in the natural abundance of isotopes, depending on their origin.)

 

Test Yourself
A sample of magnesium is found to contain 78.70% of ²⁴Mg atoms (mass 23.98 amu), 10.13% of ²⁵Mg atoms (mass 24.99 amu), and 11.17% of ²⁶Mg atoms (mass 25.98 amu). Calculate the average mass of a Mg atom.

 

Answer

24.31 amu

 

Visit this site to make mixtures of the main isotopes of the first 18 elements, gain experience with average atomic mass, and check naturally occurring isotope ratios using the Isotopes and Atomic Mass simulation.

The occurrence and natural abundances of isotopes can be experimentally determined using an instrument called a mass spectrometer. Mass spectrometry (MS) is widely used in chemistry, forensics, medicine, environmental science, and many other fields to analyze and help identify the substances in a sample of material. A mass spectrum is a plot of the relative number of ions generated versus their mass-to-charge ratios, with the height of each vertical feature or peak in a mass spectrum proportional to the fraction of particles with the specified ratio in the sample. Mass spectrometry can thus show the relative abundances of specific isotopes in a sample, directly.

See an animation that explains mass spectrometry. Watch this video from the Royal Society for Chemistry for a brief description of the rudiments of mass spectrometry.