structure of an atom essay

2.3 Atomic Structure and Symbolism

Learning objectives.

• Write and interpret symbols that depict the atomic number, mass number, and charge of an atom or ion
• Define the atomic mass unit and average atomic mass
• Calculate average atomic mass and isotopic abundance

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 2.11 ).

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, then later in terms of oxygen. Since 1961, it has been defined with regard to the most abundant isotope of carbon, atoms of which are assigned masses of exactly 12 amu. (This isotope is known as “carbon-12” as will be discussed later in this module.) Thus, one amu is exactly 1 12 1 12 of the mass of one carbon-12 atom: 1 amu = 1.6605 × × 10 −24 g. (The Dalton (Da) and the unified atomic mass unit (u) are alternative units that are equivalent to the 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 properties of these fundamental particles are summarized in Table 2.2 . (An observant student might notice that the sum of an atom’s subatomic particles does not equal the atom’s actual mass: The total mass of six protons, six neutrons, and six electrons is 12.0993 amu, slightly larger than 12.00 amu. This “missing” mass is known as the mass defect, and you will learn about it in the chapter on nuclear chemistry.)

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: A – Z = number of neutrons.

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:

Atomic charge = number of protons − number of electrons

As will be discussed in more detail, 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+). 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−).

Example 2.3

Composition of an atom.

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.

Check Your Learning

78 protons; 117 neutrons; charge is 4+

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 2.13 ). 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 2.3 . Some symbols are derived from the common name of the element; others are abbreviations of the name in another language. Most symbols have one or two letters, but three-letter symbols have been used to describe some elements that have atomic numbers greater than 112. To avoid confusion with other notations, only the first letter of a symbol is capitalized. For example, Co is the symbol for the element cobalt, but CO is the notation for the compound carbon monoxide, which contains atoms of the elements carbon (C) and oxygen (O). All known elements and their symbols are in the periodic table in [link] (also found in [link] ).

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 occasionally 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.

Link to Learning

Visit this site to learn more about IUPAC, the International Union of Pure and Applied Chemistry, and explore its periodic table.

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 2.14 ). 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 24 Mg, 25 Mg, and 26 Mg. These isotope symbols are read as “element, mass number” and can be symbolized consistent with this reading. For instance, 24 Mg is read as “magnesium 24,” and can be written as “magnesium-24” or “Mg-24.” 25 Mg 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 24 Mg atom has 12 neutrons in its nucleus, a 25 Mg atom has 13 neutrons, and a 26 Mg has 14 neutrons.

Information about the naturally occurring isotopes of elements with atomic numbers 1 through 10 is given in Table 2.4 . 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 2 H, is also called deuterium and sometimes symbolized D. Hydrogen-3, symbolized 3 H, is also called tritium and sometimes symbolized T.

Use this Build an Atom simulator to build atoms of the first 10 elements, see which isotopes exist, check nuclear stability, and gain experience with isotope symbols.

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.

For example, the element boron is composed of two isotopes: About 19.9% of all boron atoms are 10 B with a mass of 10.0129 amu, and the remaining 80.1% are 11 B with a mass of 11.0093 amu. The average atomic mass for boron is calculated to be:

It is important to understand that no single boron atom weighs exactly 10.8 amu; 10.8 amu is the average mass of all boron atoms, and individual boron atoms weigh either approximately 10 amu or 11 amu.

Example 2.4

Calculation of average atomic mass.

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.)

We can also do variations of this type of calculation, as shown in the next example.

Example 2.5

Calculation of percent abundance.

If we let x represent the fraction that is 35 Cl, then the fraction that is 37 Cl is represented by 1.00 − x .

(The fraction that is 35 Cl + the fraction that is 37 Cl must add up to 1, so the fraction of 37 Cl must equal 1.00 − the fraction of 35 Cl.)

Substituting this into the average mass equation, we have:

So solving yields: x = 0.7576, which means that 1.00 − 0.7576 = 0.2424. Therefore, chlorine consists of 75.76% 35 Cl and 24.24% 37 Cl.

69.15% Cu-63 and 30.85% Cu-65

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. In a typical mass spectrometer ( Figure 2.15 ), the sample is vaporized and exposed to a high-energy electron beam that causes the sample’s atoms (or molecules) to become electrically charged, typically by losing one or more electrons. These cations then pass through a (variable) electric or magnetic field that deflects each cation’s path to an extent that depends on both its mass and charge (similar to how the path of a large steel ball bearing rolling past a magnet is deflected to a lesser extent that that of a small steel BB). The ions are detected, and a plot of the relative number of ions generated versus their mass-to-charge ratios (a mass spectrum ) is made. The height of each vertical feature or peak in a mass spectrum is proportional to the fraction of cations with the specified mass-to-charge ratio. Since its initial use during the development of modern atomic theory, MS has evolved to become a powerful tool for chemical analysis in a wide range of applications.

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.

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Unit 2: Structure of atom

Discovery of sub-atomic particles.

• The history of atomic chemistry (Opens a modal)
• Discovery of the electron and nucleus (Opens a modal)

Atomic models

• Rutherford’s gold foil experiment (Opens a modal)
• Atomic number, mass number, and isotopes (Opens a modal)
• Isotopes (Opens a modal)
• Worked example: Identifying isotopes and ions (Opens a modal)
• Isotope composition: Counting protons, electrons, and neutrons Get 5 of 7 questions to level up!

Wave nature of electromagnetic radiation

• Light: Electromagnetic waves, the electromagnetic spectrum and photons (Opens a modal)
• Electromagnetic waves and the electromagnetic spectrum (Opens a modal)
• Introduction to light (Opens a modal)
• Spectroscopy: Interaction of light and matter (Opens a modal)
• Properties of periodic waves (Opens a modal)

Particle Nature of electromagnetic radiation: Planck's quantum theory

• Photon Energy (Opens a modal)
• Photoelectric effect (Opens a modal)

Bohr's model of hydrogen atom

• Bohr's model of hydrogen (Opens a modal)
• Absorption/emission lines (Opens a modal)
• Absorption and emission (Opens a modal)
• Emission spectrum of hydrogen (Opens a modal)
• Bohr model radii (derivation using physics) (Opens a modal)
• Bohr model radii (Opens a modal)
• Bohr model energy levels (derivation using physics) (Opens a modal)
• Bohr model energy levels (Opens a modal)
• Atomic Energy Levels (Opens a modal)

Towards Quantum mechanical model of the atom

• De Broglie wavelength (Opens a modal)
• Heisenberg uncertainty principle (Opens a modal)

Quantum mechanical model of hydrogen atom

• Quantum Wavefunction (Opens a modal)
• The quantum mechanical model of the atom (Opens a modal)
• Quantum numbers (Opens a modal)
• Quantum numbers for the first four shells (Opens a modal)
• Shells, subshells, and orbitals (Opens a modal)
• The periodic table, electron shells, and orbitals (Opens a modal)

Filling of electrons in the orbitals

• The Aufbau principle (Opens a modal)

Electronic configuration of atom

• Introduction to electron configurations (Opens a modal)
• Electron configurations article (Opens a modal)
• Noble gas configuration (Opens a modal)
• Electron configurations for the first period (Opens a modal)
• Electron configurations for the second period (Opens a modal)
• Electron configurations for the third and fourth periods (Opens a modal)
• Electron configurations of the 3d transition metals (Opens a modal)
• Electron configurations of ions (Opens a modal)
• Electron configurations Get 5 of 7 questions to level up!
• Atomic structure and electron configuration Get 3 of 4 questions to level up!

Scientific Theories of Atoms and Their Structure Essay

The development of the atomic theory.

According to Chang and Goldsby (2014), an atom is defined as, “the basic unit of an element that [participates in] chemical combinations” (p. 31). After many years of experimentation and development of theories by different scientists, it was established that an atom is made up of a positively-charged nucleus, which is surrounded by negatively-charged electrons. Moreover, different scientific theories postulated that an atom is a relatively small, indivisible particle that possesses a basic structure.

The structure of an atom includes a central, densely-packed nucleus, which is made up of smaller subatomic particles (the protons and neutrons), and it has electrons in the outer space or the “orbitals” (Chang & Goldsby, 2014; Siegfried, 2013). The current understanding of the composition of atoms and their structure has enabled scientists to develop insights into the configuration of different elements and molecules as well as the bonding behaviors and molecular polarity of various compounds (Zheng, 2012).

The history of scientific theories that explain about the composition of atoms and their structure can be traced back to the fifth century B. C. The earliest theory about atoms was made by Democritus (a Greek philosopher) who proposed that matter consisted of, “very small, indivisible particles, named atomos (meaning uncuttable or indivisible)” (Chang & Goldsby, 2014, p. 30).

However, the scientific evidence for Democritus’ claim was not available until the early 1800s when John Dalton provided a precise model of what is now known as an atom. This essay describes in greater detail the scientific theories of atoms and their structure, beginning with Dalton’s theory to the current models. In the subsequent discussions, it will become clear that there is no single scientific theory, which is made up of pure facts (Bronowski, 2011).

Instead, through scientific thinking and reasoning, scientists are capable of questioning the proposed theories before refining them in a way that brings out the truth in “hidden likenesses” (Bronowski, 2011). Moreover, through scientific theories, scientists have gained an upper hand in creating order in nature’s superficial appearances.

Dalton’s Theory

John Dalton was a scientist and schoolteacher who was the first person after Democritus to describe the chemistry behind atomism. In 1808, Dalton proposed a more precise model of an atom whereby he noted that it was a small unit of matter that could not be divided into smaller parts (Chang & Goldsby, 2014). In summary, Dalton’s theory held that an atom is the basic building block of an element.

Moreover, in any given element, Dalton noted that there were identical atoms, which had similar sizes, weight, and chemical characteristics. However, since different elements had different atoms, it was possible for the atoms of two elements to come together to form compounds (figure 1).

Nevertheless, Dalton noted that when the atoms of different elements came together to form compounds, there was no possibility that the atoms could undergo creation or destruction, but they could be separated, combined, or rearranged through a chemical reaction (Chang & Goldsby, 2014; Rocke, 2005).

In essence, Dalton’s atomic theory of matter supported Democritus’ concept that an atom could not be divided or “cut”, but it was much more detailed than the latter (Chang & Goldsby, 2014). However, from the foregoing, it is worth-noting that Dalton’s theory made no reference to the composition of atoms let alone their structure.

Despite lack of knowledge and understanding regarding the composition and the structure of atoms, Dalton noted that hydrogen atoms were different from oxygen atoms because they had dissimilar chemical properties.

Dalton’s theory was very important because it provided scientific support for the proposals made in the law of definite proportions . Additionally, the theory supported the law of multiple proportions as well as the law of conservation of mass due to its assertion that the atoms, which make up elements, cannot be created or destroyed during a chemical reaction (Rocke, 2005).

Thompson’s Atomic Model

The first attempt to describe the internal structure of an atom was made by Thompson in 1897 (Zheng, 2012). In his theory, Thompson proposed that the structure of an atom comprised a positively-charged sphere that was surrounded by negatively-charged electrons (refer to figure 2). More specifically, Thompson’s structure was known as the “plum-pudding” model because it was similar to a cake that contained raisins spread out randomly (Siegfried, 2013).

Based on this theory, Thompson took credit for discovering that an atom possessed negatively-charged particles known as electrons, which could deflect cathode rays (Zheng, 2012). Nevertheless, Thompson’s concept, particularly the deflection of cathode rays was disproved by Ernest Rutherford’s experiment in which he used alpha-particles and thin gold foils to test the deflective effect of electrons as proposed earlier (refer to figure 3).

The findings of this experiment revealed that only a small fraction of alpha-particles was deflected by the gold atoms as opposed to Thompson’s assertion that all or none of the particles would have been deflected (Zheng, 2012).

From this experiment, Rutherford proposed that an atom possesses a small centrally-placed mass of positive charge that was responsible for the sharp deflection of a fraction of the alpha-particles, which came into contact with it (Kragh, 2012). Accordingly, Rutherford’s model showed that an atom consisted of a small, dense mass of positive charge, which was called the nucleus and an empty outer space that housed the electrons (Zheng, 2012).

Bohr’s Theory

Although the aforementioned atomic theory by Rutherford was accurate in that it predicted the existence of a dense nucleus in the core of an atom, it still had a few shortcomings. For example, Rutherford proposed that the negatively-charged electrons spiraled around the dense positively-charged core. This was wrong because the laws of physics show that positive charge attracts negative charge and vice versa (Siegfried, 2013).

Therefore, if there were negative electrons in an atom, they would have been attracted to the dense positive core instead of moving around it. This led other scientists to propose that the only way that the negatively-charged electrons could be housed in an atom that contained a positively-charged nucleus was if the electrons moved around the nucleus in the same way that the planets revolve around the sun (Zheng, 2012).

This proposal was countered by Niels Bohr in his theory, which predicted that energy can be quantized. Therefore, instead of focusing on the laws of classical physics, Bohr approached the discussion on the movement of electrons around the nucleus from the perspective of quantum physics.

Accordingly, Bohr noted that the transfer of energy involved a minimum amount or a “quantum”, which is also equivalent to the energy contained in one photon. Furthermore, Bohr challenged the then existing classical laws because he believed that they did not provide sufficient evidence to explain the nature of electron movement around the nucleus.

More specifically, Bohr argued that if the classical laws were to hold, then the movement of electrons in an atom could be restricted because the electrons were in constant collision with the nucleus (Zheng, 2012). However, based on the concept of quantum physics, Bohr’s atomic theory predicted that the movement of electrons around the nucleus occurred at specific energy levels.

This means that an electron must possess a certain amount of energy to exist in a specific energy level. If an electron cannot fit in any level, then it must radiate energy until it achieves the threshold for entering a specific energy level (figure 4). Nonetheless, Bohr’s model served to explain the behavior of simple atoms such as the hydrogen atom, and failed to provide evidence for the movement of electrons within complex atoms (Siegfried, 2013; Zheng, 2012).

Modern Atomic Theories

The most current atomic models had been proposed by different scientists including Erwin Schrodinger and Werner Heisenberg. Schrodinger’s theory borrowed from the concept of the duality of matter, which had been proposed by Louis de Broglie. Schrodinger then improved Bohr’s model by predicting that the movement of electrons around the nucleus could be explained using an equation that equated electrons to waves (Zheng, 2012).

Therefore, using Schrodinger’s equation, it was possible to calculate the most probable location of an electron in the space around the nucleus. Here, note that Schrodinger’s theory differed with Bohr’s model in that it did not predict that an electron could be found in a specific location at any given time. Instead, Schrodinger proposed that there are many uncertainties surrounding the specific location of electrons in an atom.

Subsequently, Heisenberg concurred with Schrodinger’s observations by proposing that due to the duality of matter, it was not possible to determine the momentum and specific location of an electron at any given moment (Zheng, 2012). As a result, Heisenberg concluded that the momentum and location of electrons in an atom can only be determined using probability functions (refer to figure 5).

An atom refers to the smallest unit of an element, which participates in chemical reactions. During a chemical reaction, atoms cannot be created or destroyed, but instead, they can be combined, rearranged, and separated. The process of determining the properties and internal structure of an atom through the development of scientific theories can be traced back to the early nineteenth century.

Dalton’s theory is one of the well-known atomic theories that formed the basis for the development of the current understanding of atoms and their structure. Other important scientists that were involved in the development of atomic theories are Thompson, Rutherford, Bohr, and Schrodinger.

From the foregoing discussions, it appears that scientific theories are not a group of facts, but they are very important assumptions that guide scientists in the development of more refined explanations regarding the nature and existence of different things in the world.

Bronowski, J. (2011). The nature of scientific reasoning. In L. H. Peterson, J. Bizup, A. E. Fernald & M. Goldthwaite (Eds.), The Norton reader: An anthology of nonfiction (pp. 935-938). New York, NY: W. W. Norton & Company.

Chang, R., & Goldsby, K. A. (2014). General chemistry: The essential concepts (7th ed.). New York, NY: McGraw-Hill Education.

Kragh, H. (2012). Rutherford, radioactivity, and the atomic nucleus . Web.

Rocke, A. J. (2005). In search of El Dorado: John Dalton and the origins of the atomic theory. Social Research, 72 (1), 125-158.

Siegfried, T. (2013). When the atom went quantum: Bohr’s revolutionary atomic theory turns 100 . Science News. Web.

Zheng, A. (2012). The evolution of atomic theory. Young Scientists Journal, 1 (12), 74 76. Web.

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Home — Essay Samples — Science — Atom — The Structure of an Atom

The Structure of an Atom

• Categories: Atom

About this sample

Words: 454 |

Published: Jan 15, 2019

Words: 454 | Page: 1 | 3 min read

Works Cited

• Chang, R. (2010). Chemistry (10th ed.). McGraw Hill.
• Housecroft, C. E., & Sharpe, A. G. (2018). Inorganic Chemistry (5th ed.). Pearson.
• Huheey, J. E., Keiter, E. A., & Keiter, R. L. (2014). Inorganic chemistry: principles of structure and reactivity (4th ed.). Pearson.
• Kotz, J. C., Treichel Jr, P. M., & Townsend, J. R. (2016). Chemistry and Chemical Reactivity (9th ed.). Cengage Learning.
• Martin, G. J., & Cockett, M. C. R. (2000). Essential Chemistry for Cambridge IGCSE (2nd ed.). Oxford University Press.
• McMurry, J., & Fay, R. C. (2017). Chemistry (7th ed.). Pearson.
• Moore, J. W., & Stanitski, C. L. (2017). Chemistry: The Molecular Science (5th ed.). Cengage Learning.
• Petrucci, R. H., Herring, F. G., Madura, J. D., & Bissonnette, C. (2017). General Chemistry: Principles and Modern Applications (11th ed.). Pearson.
• Silberberg, M. S. (2016). Chemistry: The Molecular Nature of Matter and Change (8th ed.). McGraw Hill Education.
• Zumdahl, S. S., & DeCoste, D. J. (2016). Chemistry (10th ed.). Cengage Learning.

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• Atomic Structure

Atomic Structure - Discovery of Subatomic Particles

The atomic structure refers to the structure of an atom comprising a nucleus (centre) in which the protons (positively charged) and neutrons (neutral) are present. The negatively charged particles called electrons revolve around the centre of the nucleus .

Download Complete Chapter Notes of Structure of Atom Download Now

The history of atomic structure and quantum mechanics dates back to the times of Democritus, the person who first proposed that matter is composed of atoms. The study of the structure of an atom gives a great insight into the entire class of chemical reactions, bonds and their physical properties. The first scientific theory of atomic structure was proposed by John Dalton in the 1800s.

Atomic Structure Quick Revision for the JEE

Structure of Atom – Important Topics

Table of Contents

What is atomic structure, atomic models, dalton’s atomic theory, thomson atomic model, rutherford atomic theory, subatomic particles, atomic structure of isotopes, bohr’s atomic theory, dual nature of matter.

The advances in atomic structure and quantum mechanics have led to the discovery of other fundamental particles. The discovery of subatomic particles has been the base for many other discoveries and inventions.

The atomic structure of an element refers to the constitution of its nucleus and the arrangement of the electrons around it. Primarily, the atomic structure of matter is made up of protons , electrons and neutrons.

The protons and neutrons make up the nucleus of the atom, which is surrounded by the electrons belonging to the atom. The atomic number of an element describes the total number of protons in its nucleus.

Neutral atoms have equal numbers of protons and electrons. However, atoms may gain or lose electrons in order to increase their stability, and the resulting charged entity is called an ion.

Atoms of different elements have different atomic structures because they contain different numbers of protons and electrons . This is the reason for the unique characteristics of different elements.

In the 18th and 19th centuries, many scientists attempted to explain the structure of the atom with the help of atomic models. Each of these models had its own merits and demerits and was pivotal to the development of the modern atomic model . The most notable contributions to the field were by the scientists such as John Dalton, J.J. Thomson, Ernest Rutherford and Niels Bohr. Their ideas on the structure of the atom are discussed in this subsection.

The English chemist John Dalton suggested that all matter is made up of atoms, which were indivisible and indestructible. He also stated that all the atoms of an element were exactly the same, but the atoms of different elements differ in size and mass.

Chemical reactions, according to Dalton’s atomic theory, involve a rearrangement of atoms to form products. According to the postulates proposed by Dalton, the atomic structure comprises atoms, the smallest particle responsible for the chemical reactions to occur.

The following are the postulates of his theory:

• Every matter is made up of atoms.
• Atoms are indivisible.
• Specific elements have only one type of atom in them.
• Each atom has its own constant mass that varies from element to element.
• Atoms undergo rearrangement during a chemical reaction.
• Atoms can neither be created nor destroyed but can be transformed from one form to another.

Dalton’s atomic theory successfully explained the Laws of chemical reactions , namely, the Law of conservation of mass, the Law of constant properties, the Law of multiple proportions and the Law of reciprocal proportions.

Demerits of Dalton’s Atomic Theory

• The theory was unable to explain the existence of isotopes.
• Nothing about the structure of the atom was appropriately explained.
• Later, scientists discovered particles inside the atom that proved the atoms are divisible.

The discovery of particles inside atoms led to a better understanding of chemical species; these particles inside the atoms are called subatomic particles. The discovery of various subatomic particles is as follows:

The English chemist Sir Joseph John Thomson put forth his model describing the atomic structure in the early 1900s.

He was later awarded the Nobel Prize for the discovery of “electrons” . His work is based on an experiment called the cathode ray experiment . The construction of working of the experiment is as follows:

Cathode Ray Experiment

It has a tube made of glass which has two openings, one for the vacuum pump and the other for the inlet through which a gas is pumped in.

The role of the vacuum pump is to maintain a “partial vacuum” inside the glass chamber. A high-voltage power supply is connected using electrodes, i.e., cathode and anode , which are fitted inside the glass tube.

Observations:

• When a high voltage power supply is switched on, there are rays emerging from the cathode towards the anode. This was confirmed by the ‘Fluorescent spots’ on the ZnS screen used. These rays were called “Cathode Rays”.
• When an external electric field is applied, the cathode rays get deflected towards the positive electrode, but in the absence of an electric field, they travel in a straight line.

• With all this evidence, Thompson concluded that cathode rays are made of negatively charged particles called “electrons”.
• On applying the electric and magnetic field upon the cathode rays (electrons), Thomson found the charge-to-mass ratio (e/m) of electrons. (e/m) for electron: 17588 × 10 11 e/bg.

From this ratio, the charge of the electron was found by Mullikin through an oil drop experiment . [Charge of e – = 1.6 × 10 -16 C and Mass of e – = 9.1093 × 10 -31 kg].

Conclusions:

Based on conclusions from his cathode ray experiment, Thomson described the atomic structure as a positively charged sphere into which negatively charged electrons were embedded.

It is commonly referred to as the “plum pudding model” because it can be visualised as a plum pudding dish where the pudding describes the positively charged atom and the plum pieces describe the electrons.

Thomson’s atomic structure described atoms as electrically neutral, i.e., the positive and the negative charges were of equal magnitude.

Limitations of Thomson’s Atomic Structure:  Thomson’s atomic model does not clearly explain the stability of an atom. Also, further discoveries of other subatomic particles couldn’t be placed inside his atomic model.

Rutherford, a student of J. J. Thomson, modified the atomic structure with the discovery of another subatomic particle called “Nucleus” . His atomic model is based on the Alpha ray scattering experiment.

Alpha Ray Scattering Experiment

Construction:.

• A very thin gold foil of 1000 atoms thick is taken.
• Alpha rays (doubly charged Helium He 2+ ) were made to bombard the gold foil.
• Zn S screen is placed behind the gold foil.
• Most of the rays just went through the gold foil, making scintillations (bright spots) in the ZnS screen.
• A few rays got reflected after hitting the gold foil.
• One in 1000 rays got reflected by an angle of 180° (retraced path) after hitting the gold foil.
• Since most rays passed through, Rutherford concluded that most of the space inside the atom is empty.
• A few rays got reflected because of the repulsion of its positive with some other positive charge inside the atom.
• 1/1000th of the rays got strongly deflected because of a very strong positive charge in the centre of the atom. He called this strong positive charge “nucleus”.
• He said most of the charge and mass of the atom resides in the nucleus.

Rutherford’s Structure of Atom

Based on the above observations and conclusions, Rutherford proposed his own atomic structure , which is as follows.

• The nucleus is at the centre of an atom, where most of the charge and mass is concentrated.
• The atomic structure is spherical.
• Electrons revolve around the nucleus in a circular orbit, similar to the way planets orbit the sun.

Limitations of the Rutherford Atomic Model

• If electrons have to revolve around the nucleus, they will spend energy and that too against the strong force of attraction from the nucleus, a lot of energy will be spent by the electrons, and eventually, they will lose all their energy and will fall into the nucleus so the stability of atom is not explained.
• If electrons continuously revolve around the ‘nucleus, the type of spectrum expected is a continuous spectrum. But in reality, what we see is a line spectrum.

Atomic Structure – Rutherford’s Model, J.J Thomson’s Model

• Protons are positively charged subatomic particles. The charge of a proton is 1e, which corresponds to approximately 1.602 × 10 -19
• The mass of a proton is approximately 1.672 × 10 -24
• Protons are over 1800 times heavier than electrons.
• The total number of protons in the atoms of an element is always equal to the atomic number of the element.
• The mass of a neutron is almost the same as that of a proton, i.e., 1.674×10 -24
• Neutrons are electrically neutral particles and carry no charge.
• Different isotopes of an element have the same number of protons but vary in the number of neutrons present in their respective nuclei.
• The charge of an electron is -1e, which approximates to -1.602 × 10 -19
• The mass of an electron is approximately 9.1 × 10 -31 .
• Due to the relatively negligible mass of electrons, they are ignored when calculating the mass of an atom.

Nucleons are the components of the nucleus of an atom. A nucleon can either be a proton or a neutron. Each element has a unique number of protons in it, which is described by its unique atomic number . However, several atomic structures of an element can exist, which differ in the total number of nucleons.

These variants of elements having a different nucleon number (also known as the mass number) are called isotopes of the element. Therefore, the isotopes of an element have the same number of protons but differ in the number of neutrons.

The atomic structure of an isotope is described with the help of the chemical symbol of the element, the atomic number of the element and the mass number of the isotope. For example, there exist three known naturally occurring isotopes of hydrogen , namely, protium, deuterium and tritium. The atomic structures of these hydrogen isotopes are illustrated below.

The isotopes of an element vary in stability. The half-lives of isotopes also differ. However, they generally have similar chemical behaviour owing to the fact that they hold the same electronic structures .

Atomic Structures of Some Elements

The structure of an atom of an element can be simply represented via the total number of protons, electrons and neutrons present in it. The atomic structures of a few elements are illustrated below.

The most abundant isotope of hydrogen on the planet Earth is protium. The atomic number and the mass number of this isotope are 1 and 1, respectively.

Structure of Hydrogen Atom: This implies that it contains one proton, one electron and no neutrons (Total number of neutrons = Mass number – Atomic number)

Carbon has two stable isotopes – 12C and 13C. Of these isotopes, 12C has an abundance of 98.9%. It contains 6 protons, 6 electrons and 6 neutrons.

Structure of Carbon Atom: The electrons are distributed into two shells, and the outermost shell (valence shell) has four electrons. The tetravalency of carbon enables it to form a variety of chemical bonds with various elements.

There exist three stable isotopes of oxygen – 18O, 17O and 16O. However, oxygen-16 is the most abundant isotope.

Structure of Oxygen Atom: Since the atomic number of this isotope is 8 and the mass number is 16, it consists of 8 protons and 8 neutrons. 6 out of the 8 electrons in an oxygen atom lie in the valence shell.

Neils Bohr put forth his model of the atom in the year 1915. This is the most widely used atomic model to describe the atomic structure of an element which is based on Planck’s theory of quantization .

Postulates:

• The electrons inside atoms are placed in discrete orbits called “stationery orbits”.
• The energy levels of these shells can be represented via quantum numbers.
• Electrons can jump to higher levels by absorbing energy and move to lower energy levels by losing or emitting their energy.
• As long as an electron stays in its own stationery, there will be no absorption or emission of energy.
• Electrons revolve around the nucleus in these stationary orbits only.
• The energy of the stationary orbits is quantised.

Limitations of Bohr’s Atomic Theory:

• Bohr’s atomic structure works only for single electron species such as H, He+, Li2+, Be3+, ….
• When the emission spectrum of hydrogen was observed under a more accurate spectrometer, each line spectrum was seen to be a combination of a number of smaller discrete lines.
• Both Stark and Zeeman’s effects couldn’t be explained using Bohr’s theory.

Heisenberg’s uncertainty principle: Heisenberg stated that no two conjugate physical quantities could be measured simultaneously with 100% accuracy. There will always be some error or uncertainty in the measurement.

Drawback: Position and momentum are two such conjugate quantities that were measured accurately by Bohr (theoretically).

Stark effect: Phenomenon of deflection of electrons in the presence of an electric field.

Zeeman effect: Phenomenon of deflection of electrons in the presence of a magnetic field.

The electrons, which were treated to be particles, and the evidence of the photoelectric effect show they also have a wave nature. This was proved by Thomas Young with the help of his double-slit experiment .

De-Broglie concluded that since nature is symmetrical, so should light or any other matter wave be.

Quantum Numbers

• Principal Quantum Number (n): It denotes the orbital number or shell number of an electron.
• Azimuthal Quantum Numbers ( l ): It denotes the orbital (sub-orbit) of the electron.
• Magnetic Quantum Number: It denotes the number of energy states in each orbit.
• Spin Quantum number(s): It denotes the direction of spin, S = -½ = Anticlockwise and ½ = Clockwise.

Electronic Configuration of an Atom

The electrons have to be filled in the s, p, d and f in accordance with the following rule.

1. Aufbau’s principle: The filling of electrons should take place in accordance with the ascending order of energy of orbitals.

• Lower energy orbital should be filled first, and higher energy levels.
• The energy of orbital α(p + l) value it two orbitals have the same (n + l ) value, E α n
• Ascending order of energy 1s, 2s, 2p, 3s, 3p, 4s, 3d, . . .

2. Pauli’s exclusion principle: No two electrons can have all four quantum numbers to be the same, or if two electrons have to be placed in an energy state, they should be placed with opposite spies.

3. Hund’s rule of maximum multiplicity: In the case of filling degenerate (same energy) orbitals, all the degenerate orbitals have to be singly filled first, and then, only pairing has to happen.

Atomic Structure Solved Problems and Solutions

Atomic Structure – Important Questions

Structure of Atom Class 11 – Full Chapter Revision

Structure of Atom – Top 12 Most Important JEE Main Questions

Frequently Asked Questions on Atomic Structure

What are subatomic particles.

Subatomic particles are the particles that constitute an atom. Generally, this term refers to protons, electrons and neutrons.

How do the atomic structures of isotopes vary?

They vary in terms of the total number of neutrons present in the nucleus of the atom, which is described by their nucleon numbers.

What are the shortcomings of Bohr’s atomic model?

According to this atomic model, the structure of an atom offers poor spectral predictions for larger atoms. It also failed to explain the Zeeman effect. It could only successfully explain the hydrogen spectrum.

How can the total number of neutrons in the nucleus of a given isotope be determined?

The mass number of an isotope is given by the sum of the total number of protons and neutrons in it. The atomic number describes the total number of protons in the nucleus. Therefore, the number of neutrons can be determined by subtracting the atomic number from the mass number.

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Introduction to Structure of Atom

An atom is present at the most basic level in everything we see around us. In fact, atoms compose every living organism . Matter makes up every non-living thing around us such as tables, chairs, water, etc. But the building blocks of matter are atoms. Thus, the composition of everything, living or non-living, is atoms. Let us study about atom and structure of atom.

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Atom is a Greek word which means “ indivisible .” The Greeks believed that matter can be broken down into very small invisible particles called atoms. Greek philosophers such as Democritus and John Dalton put forward the concept of the atom.

Democritus explained the nature of matter. He proposed that  matter  makes up all substances. Furthermore, he stated that atoms are constantly moving. They are invisible, minuscule particles that are different in shape, size, and  temperature . And that we cannot destroy atoms.

Learn the concept of the Atomic number here in detail .

Later in the year 1808, John Dalton proposed the atomic theory. He explained the law of chemical combination . By the end of 18 th and the early 20 th centuries, many scientists developed and proposed several concepts on “atom.” Namely, J.J Thomson , Gold stein, Rutherford , and  Bohr among others.

Atom is the smallest unit of matter. It consists of a positively charged center termed as “ nucleus ”. Also, negatively charged electrons  surround the central nucleus. Even though an atom is the smallest unit of matter it retains all the chemical properties of an element . For example, silver atoms and a few other constituents make up a silver spoon. The tiny subatomic particles that compose a silver atom define the properties of the silver atom.

Atoms are further arranged and organized to form larger structures known as molecules . Atoms and molecules adhere to the general chemistry and physics rule even when they are part of living human body. Now let us study the structure of atom to further understand how the atoms react, behave and interact.

Learn more about  Thomson’s Model of Atom in detail .

Structure of Atom

The structure of atom consists of two parts:

• an atomic nucleus
• extra nucleus part

The tiny atomic nucleus is the center of an atom. It constitutes positively charged particles “protons” and uncharged particles “ neutrons .” On the other hand, the extra nucleus part is a much larger region. It consists of a cloud of negatively charged particles called an electron.  Electrons   revolve in orbit  around the nucleus. The attraction between the protons and electrons holds the structure of atom together.

Learn how electrons are distributed in different shells in detail here .

Generally, all atoms except hydrogen consist of these three subatomic particles. Hydrogen is an exception to all atoms as it contains just one proton and one electron but lacks neutrons. The number of protons indicates what element an atom is.  Whereas the number of electrons indicates the type of reactions that will happen in an atom.

The atomic nucleus in the structure of atom consists of a fixed number of protons. Also, the proton attracts the same number of electrons thereby making an atom electrically neutral. The addition or removal of electrons from an atom results in the formation of ions .

Learn about the Disadvantages of Rutherford’s Atomic Model here .

You can download Structure of Atom Cheat Sheet by clicking on the download button below

Discovery of an Electron

In the year 1897, a British Physicist named J.J Thompson  proposed that an atom constitutes of at least one negatively charged particle. He named it “corpuscles” which was later called “electron.”

‘e’ represents an electron and it contributes to the negative charge of an atom. The absolute charge of an electron is the negative charge of 1.6×10 -19 coulombs. The relative mass of an electron is 1/1836. Furthermore, the mass of an electron is 0.

Discovery of Proton

Proton was discovered by Rutherford when he conducted the famous gold foil experiment. In 1886 Goldstein discovered the presence of positively charged rays while experimenting with the discharged tube using perforated cathode. The rays were named as anode rays or canal rays. A series of experiments led to the discovery of protons. Protons are the particles that contribute to the positive charge of the atom.

“p” represents proton. The absolute charge of a proton is the positive charge of 1.6×10 -19 coulomb . The mass of a proton is 1.6×10 -24 g and is considered 1 that is mass of a hydrogen atom .

Discovery of Neutron

The discovery of neutron didn’t happen until the year 1932. James Chadwick discovered the neutron. He used scattered particle to calculate the mass of the neutral particle. The subatomic particle “neutron” is present in an atom’s nucleus.“n” represents neutron. It is a neutral particle. The mass of a neutron is 1.6 x 10 -24 g.

Gram is not an appropriate unit for the calculation of such tiny subatomic particles . But Dalton or amu (atomic mass unit) is appropriate. Furthermore, protons and neutrons have a mass that is nearly 1 amu.

Learn more about Bohr’s Atomic Model here .

Solved Question for You

Question: What is the net charge of an atom?

Ans: There is no net charge of an atom. The charge of electrons is negative whereas the charge of protons is positive. The equal positive charge of the proton and the negative charge of the electron cancel each other. Therefore, the atom has no net charge. In a neutral atom, the number of electrons revolving around the nucleus and the number of protons inside the nucleus are equal in number.

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2.5 The Structure of The Atom

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  Learning Objectives

• Write and interpret symbols that depict the atomic number, mass number, and charge of an atom or ion
• Define the atomic mass unit and average atomic mass
• Calculate average atomic mass and isotopic abundance

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 \(\PageIndex{1}\)).

Atoms—and the protons, neutrons, and electrons that compose them—are extremely small. For example, a carbon atom weighs less than 2 \(\times\) 10 −23 g, and an electron has a charge of less than 2 \(\times\) 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, then later in terms of oxygen. Since 1961, it has been defined with regard to the most abundant isotope of carbon, atoms of which are assigned masses of exactly 12 amu. (This isotope is known as “carbon-12” as will be discussed later in this module.) Thus, one amu is exactly \(1/12\) of the mass of one carbon-12 atom: 1 amu = 1.6605 \(\times\) 10 −24 g. (The Dalton (Da) and the unified atomic mass unit (u) are alternative units that are equivalent to the 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 \(\times\) 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 properties of these fundamental particles are summarized in Table \(\PageIndex{1}\). (An observant student might notice that the sum of an atom’s subatomic particles does not equal the atom’s actual mass: The total mass of six protons, six neutrons, and six electrons is 12.0993 amu, slightly larger than the 12.00 amu of an actual carbon-12 atom. This “missing” mass is known as the mass defect, and you will learn about it in the chapter on nuclear chemistry.)

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: A – Z = number of neutrons.

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

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:

Atomic charge = number of protons − 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+). 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−).

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 (Figure \(\PageIndex{2}\)).

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.

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+

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 \(\PageIndex{3}\)). 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 \(\PageIndex{2}\). 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 Cl for chlorine. To avoid confusion with other notations, only the first letter of a symbol is capitalized. For example, Co is the symbol for the element cobalt, but CO is the notation for the compound carbon monoxide, which contains atoms of the elements carbon (C) and oxygen (O). All known elements and their symbols are in the periodic table.

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.

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 \(\PageIndex{4}\)). 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 24 Mg, 25 Mg, and 26 Mg. These isotope symbols are read as “element, mass number” and can be symbolized consistent with this reading. For instance, 24 Mg is read as “magnesium 24,” and can be written as “magnesium-24” or “Mg-24.” 25 Mg 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 24 Mg atom has 12 neutrons in its nucleus, a 25 Mg atom has 13 neutrons, and a 26 Mg has 14 neutrons.

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 2 H, is also called deuterium and sometimes symbolized D. Hydrogen-3, symbolized 3 H, is also called tritium and sometimes symbolized T.

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.

\[\mathrm{average\: mass}=\sum_{i}(\mathrm{fractional\: abundance\times isotopic\: mass})_i \nonumber \]

For example, the element boron is composed of two isotopes: About 19.9% of all boron atoms are 10 B with a mass of 10.0129 amu, and the remaining 80.1% are 11 B with a mass of 11.0093 amu. The average atomic mass for boron is calculated to be:

\[\begin{align*} \textrm{boron average mass} &=\mathrm{(0.199\times10.0129\: amu)+(0.801\times11.0093\: amu)}\\ &=\mathrm{1.99\: amu+8.82\: amu}\\ &=\mathrm{10.81\: amu} \end{align*} \nonumber \]

It is important to understand that no single boron atom weighs exactly 10.8 amu; 10.8 amu is the average mass of all boron atoms, and individual boron atoms weigh either approximately 10 amu or 11 amu.

Example \(\PageIndex{2}\): Calculation of Average Atomic Mass

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% 20 Ne (mass 19.9924 amu), 0.47% 21 Ne (mass 20.9940 amu), and 7.69% 22 Ne (mass 21.9914 amu). What is the average mass of the neon in the solar wind?

\[\begin{align*} \mathrm{average\: mass} &=\mathrm{(0.9184\times19.9924\: amu)+(0.0047\times20.9940\: amu)+(0.0769\times21.9914\: amu)}\\ &=\mathrm{(18.36+0.099+1.69)\:amu}\\ &=\mathrm{20.15\: amu} \end{align*} \nonumber \]

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.)

Exercise \(\PageIndex{2}\)

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

We can also do variations of this type of calculation, as shown in the next example.

Example \(\PageIndex{3}\): Calculation of Percent Abundance

Naturally occurring chlorine consists of 35 Cl (mass 34.96885 amu) and 37 Cl (mass 36.96590 amu), with an average mass of 35.453 amu. What is the percent composition of Cl in terms of these two isotopes?

The average mass of chlorine is the fraction that is 35 Cl times the mass of 35 Cl plus the fraction that is 37 Cl times the mass of 37 Cl.

\[\mathrm{average\: mass=(fraction\: of\: ^{35}Cl\times mass\: of\: ^{35}Cl)+(fraction\: of\: ^{37}Cl\times mass\: of\: ^{37}Cl)} \nonumber \]

If we let x represent the fraction that is 35 Cl, then the fraction that is 37 Cl is represented by 1.00 − x .

(The fraction that is 35 Cl + the fraction that is 37 Cl must add up to 1, so the fraction of 37 Cl must equal 1.00 − the fraction of 35 Cl.)

Substituting this into the average mass equation, we have:

\[\begin{align*} \mathrm{35.453\: amu} &=(x\times 34.96885\: \ce{amu})+[(1.00-x)\times 36.96590\: \ce{amu}]\\ 35.453 &=34.96885x+36.96590-36.96590x\\ 1.99705x &=1.513\\ x&=\dfrac{1.513}{1.99705}=0.7576 \end{align*} \nonumber \]

So solving yields: x = 0.7576, which means that 1.00 − 0.7576 = 0.2424. Therefore, chlorine consists of 75.76% 35 Cl and 24.24% 37 Cl.

Exercise \(\PageIndex{3}\)

Naturally occurring copper consists of 63 Cu (mass 62.9296 amu) and 65 Cu (mass 64.9278 amu), with an average mass of 63.546 amu. What is the percent composition of Cu in terms of these two isotopes?

69.15% Cu-63 and 30.85% Cu-65

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. In a typical mass spectrometer (Figure \(\PageIndex{5}\)), the sample is vaporized and exposed to a high-energy electron beam that causes the sample’s atoms (or molecules) to become electrically charged, typically by losing one or more electrons. These cations then pass through a (variable) electric or magnetic field that deflects each cation’s path to an extent that depends on both its mass and charge (similar to how the path of a large steel ball bearing rolling past a magnet is deflected to a lesser extent that that of a small steel BB). The ions are detected, and a plot of the relative number of ions generated versus their mass-to-charge ratios (a mass spectrum ) is made. The height of each vertical feature or peak in a mass spectrum is proportional to the fraction of cations with the specified mass-to-charge ratio. Since its initial use during the development of modern atomic theory,   MS has evolved to become a powerful tool for chemical analysis in a wide range of applications.

Video \(\PageIndex{1}\) : Watch this video from the Royal Society for Chemistry for a brief description of the rudiments of mass spectrometry.

An atom consists of a small, positively charged nucleus surrounded by electrons. The nucleus contains protons and neutrons; its diameter is about 100,000 times smaller than that of the atom. The mass of one atom is usually expressed in atomic mass units (amu), which is referred to as the atomic mass. An amu is defined as exactly \(1/12\) of the mass of a carbon-12 atom and is equal to 1.6605 \(\times\) 10 −24 g.

Protons are relatively heavy particles with a charge of 1+ and a mass of 1.0073 amu. Neutrons are relatively heavy particles with no charge and a mass of 1.0087 amu. Electrons are light particles with a charge of 1− and a mass of 0.00055 amu. The number of protons in the nucleus is called the atomic number (Z) and is the property that defines an atom’s elemental identity. The sum of the numbers of protons and neutrons in the nucleus is called the mass number and, expressed in amu, is approximately equal to the mass of the atom. An atom is neutral when it contains equal numbers of electrons and protons.

Isotopes of an element are atoms with the same atomic number but different mass numbers; isotopes of an element, therefore, differ from each other only in the number of neutrons within the nucleus. When a naturally occurring element is composed of several isotopes, the atomic mass of the element represents the average of the masses of the isotopes involved. A chemical symbol identifies the atoms in a substance using symbols, which are one-, two-, or three-letter abbreviations for the atoms.

Key Equations

• \(\mathrm{average\: mass}=\sum_{i}(\mathrm{fractional\: abundance \times isotopic\: mass})_i\)

Paul Flowers (University of North Carolina - Pembroke), Klaus Theopold (University of Delaware) and Richard Langley (Stephen F. Austin State University) with contributing authors.  Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution License 4.0 license. Download for free at http://cnx.org/contents/[email protected] ).

Essay on Atoms

Students are often asked to write an essay on Atoms in their schools and colleges. And if you’re also looking for the same, we have created 100-word, 250-word, and 500-word essays on the topic.

Let’s take a look…

100 Words Essay on Atoms

What are atoms.

Atoms are the tiny building blocks of everything around us. They are so small that we can’t see them with our eyes. They make up the air we breathe, the food we eat, and even us!

Structure of an Atom

An atom is made up of three smaller parts: protons, neutrons, and electrons. Protons and neutrons are in the middle, called the nucleus. Electrons move around the nucleus in paths called orbits.

Atomic Number and Mass

Every atom has a unique number of protons, called its atomic number. The total number of protons and neutrons gives the atomic mass.

Chemical Reactions and Atoms

In chemical reactions, atoms join together or split apart. They can form molecules, which are groups of atoms. For example, two hydrogen atoms and one oxygen atom make a water molecule.

Sometimes, atoms of the same element have different numbers of neutrons. These are called isotopes. For example, carbon-12 and carbon-14 are isotopes of carbon.

Atoms and Energy

Atoms can also store and release energy. This is used in nuclear power plants and atomic bombs. But don’t worry, atoms in our daily life are safe and won’t explode!

250 Words Essay on Atoms

Atoms are tiny particles that make up everything around us. They are so small that we cannot see them with our eyes or even with a normal microscope. Atoms are the building blocks of matter, which means they are the basis for everything in the universe.

Parts of an Atom

An atom is made up of three types of particles: protons, neutrons, and electrons. Protons and neutrons are in the center of the atom, called the nucleus. Protons have a positive charge, and neutrons have no charge. Electrons, which have a negative charge, move around the nucleus.

Atoms and Elements

Atoms make up elements. Elements are pure substances that consist of only one type of atom. For example, a gold ring is made up of gold atoms, and a silver spoon is made up of silver atoms. There are about 118 known elements, and they are listed in the Periodic Table.

Importance of Atoms

Atoms are important because they form the basis of everything. They make up the air we breathe, the food we eat, and even our bodies. Scientists study atoms to understand how matter works and to develop new materials and technologies.

In conclusion, atoms are tiny but mighty. They may be small, but they play a big role in everything we do. Understanding atoms helps us understand the world around us.

500 Words Essay on Atoms

Atoms are tiny particles that make up everything in the universe. Imagine a tiny dot that is so small that it is impossible to see with your eyes. That’s how small an atom is! It’s like the building block of all matter. Matter refers to anything you can touch, see, or feel. So, whether it’s a dog, a tree, or even a book, they are all made up of atoms.

An atom is made up of three smaller parts. These parts are called protons, neutrons, and electrons. Protons and neutrons stay in the center of the atom, which is called the nucleus. Protons have a positive charge and neutrons have no charge. Electrons, which are negatively charged, move around the nucleus in paths called orbits.

Types of Atoms

Different types of atoms are called elements. You might have seen the periodic table in your science class. That table lists all the known types of atoms or elements. There are 118 elements that we know of. Some of them are very common like hydrogen, oxygen, and carbon. Others are rare like gold and platinum.

Atoms and Chemical Reactions

Atoms can join together to form molecules. For example, two hydrogen atoms and one oxygen atom can join together to form a water molecule. This process is called a chemical reaction. Chemical reactions are happening all around us and inside us. For example, our body uses chemical reactions to break down food and get energy.

Atoms also have a lot to do with energy. The sun, which gives us light and warmth, is a big ball of atoms undergoing a reaction called nuclear fusion. In this process, atoms combine to form a larger atom and release a lot of energy. On the other hand, nuclear power plants use a process called nuclear fission where a large atom is split into smaller atoms, releasing energy.

Atoms are very important because they make up everything in the universe. Understanding atoms helps us understand the world around us. Scientists study atoms to find new ways to create energy, make new materials, and even cure diseases.

In conclusion, atoms may be tiny, but they are mighty. They make up everything we see, touch, and feel. They are involved in chemical reactions that are essential for life. They are also a key part of how we get energy. So, even though we can’t see them, atoms are everywhere and they are very important!

That’s it! I hope the essay helped you.

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