5 hypothesis of dalton's atomic theory

If you're seeing this message, it means we're having trouble loading external resources on our website.

If you're behind a web filter, please make sure that the domains *.kastatic.org and *.kasandbox.org are unblocked.

To log in and use all the features of Khan Academy, please enable JavaScript in your browser.

Chemistry library

Course: chemistry library   >   unit 5.

• The history of atomic chemistry

Dalton's atomic theory

• Discovery of the electron and nucleus
• Rutherford’s gold foil experiment
• Bohr's model of hydrogen
• Dalton's atomic theory was the first complete attempt to describe all matter in terms of atoms and their properties.
• Dalton based his theory on the law of conservation of mass and the law of constant composition .
• The first part of his theory states that all matter is made of atoms, which are indivisible .
• The second part of the theory says all atoms of a given element are identical in mass and properties .
• The third part says compounds are combinations of two or more different types of atoms .
• The fourth part of the theory states that a chemical reaction is a rearrangement of atoms .
• Parts of the theory had to be modified based on the discovery of subatomic particles and isotopes.

Chemists ask questions.

Basis for dalton's theory, part 1: all matter is made of atoms., part 2: all atoms of a given element are identical in mass and properties., part 3: compounds are combinations of two or more different types of atoms., part 4: a chemical reaction is a rearrangement of atoms., what have we learned since dalton proposed his theory, attributions:.

• OpenStax College. "Early Ideas in Atomic Theory." OpenStax CNX. October 2, 2014. http://cnx.org/contents/[email protected]:HdZmYjzP@4/Early-Ideas-in-Atomic-Theory . CC-BY 4.0
• "Atomic Theory." UC Davis ChemWiki. http://chemwiki.ucdavis.edu/Physical_Chemistry/Atomic_Theory/Atomic_Theory . CC-BY-NC-SA 3.0 US
• The first part of his theory states that all matter is made of atoms, which are indivisible.
• The second part of the theory says all atoms of a given element are identical in mass and properties.
• The third part says compounds are combinations of two or more different types of atoms.
• The fourth part of the theory states that a chemical reaction is a rearrangement of atoms.
• Parts of the theory had to be modified based on the existence of subatomic particles and isotopes.

Want to join the conversation?

• Upvote Button navigates to signup page
• Downvote Button navigates to signup page
• Flag Button navigates to signup page

• Structure of Atom
• Daltons Atomic Theory

Dalton’s Atomic Theory

What is atomic theory.

Dalton’s atomic theory was a scientific theory on the nature of matter put forward by the English physicist and chemist John Dalton in the year 1808. It stated that all matter was made up of small, indivisible particles known as ‘atoms’.

All substances, according to Dalton’s atomic theory, are made up of atoms, which are indivisible and indestructible building units. While an element’s atoms were all the same size and mass, various elements possessed atoms of varying sizes and masses.

Table of Contents

Postulates of dalton’s atomic theory, limitations of dalton’s atomic theory, what are the merits of dalton’s atomic theory, recommended videos.

• Frequently Asked Questions – FAQs
• All matter is made up of tiny, indivisible particles called atoms.
• All atoms of a specific element are identical in mass, size, and other properties. However, atoms of different element exhibit different properties and vary in mass and size.
• Atoms can neither be created nor destroyed. Furthermore, atoms cannot be divided into smaller particles.
• Atoms of different elements can combine with each other in fixed whole-number ratios in order to form compounds.
• Atoms can be rearranged, combined, or separated in chemical reactions.
• It does not account for subatomic particles: Dalton’s atomic theory stated that atoms were indivisible. However, the discovery of subatomic particles (such as protons, electrons, and neutrons) disproved this postulate.
• It does not account for isotopes: As per Dalton’s atomic theory, all atoms of an element have identical masses and densities. However, different isotopes of elements have different atomic masses (Example: hydrogen, deuterium, and tritium).
• It does not account for isobars: This theory states that the masses of the atoms of two different elements must differ. However, it is possible for two different elements to share the same mass number. Such atoms are called isobars (Example: 40 Ar and 40 Ca).
• Elements need not combine in simple, whole-number ratios to form compounds: Certain complex organic compounds do not feature simple ratios of constituent atoms. Example: sugar/sucrose (C 11 H 22 O 11 ).
• The theory does not account for allotropes: The differences in the properties of diamond and graphite, both of which contain only carbon, cannot be explained by Dalton’s atomic theory.

Dalton’s Atomic Theory – The Indestructible Atoms

• The law of multiple proportions, the law of conservation of mass, and the law of constant proportions are not violated by Dalton’s atomic theory.
• The theory provides a basis to differentiate between elements and compounds.

The Periodic Table

Atomic Structure

To learn more about Dalton’s atomic theory and other related concepts such as Thomson’s atomic model , register with BYJU’S and download the mobile application on your smartphone.

Frequently Asked Questions – FAQs

How does dalton’s atomic theory explain the law of conservation of mass.

Since it states that atoms cannot be created or destroyed, Dalton’s theory suggests that the net mass of the participating species in a chemical reaction is conserved. This postulate, therefore, accounts for the law of conservation of mass.

How does Dalton’s atomic theory differentiate between elements and compounds?

This theory states that elements combine in fixed, whole-number ratios to form compounds. Therefore, it suggests that compounds are made up of molecules that contain two or more atoms of different elements.

What are the 5 key postulates of Dalton’s atomic theory?

The 5 postulates of Dalton’s atomic theory are listed below.

• All matter is made up of atoms, which are tiny, indivisible particles.
• All the atoms of an element have the same size, mass, and properties but the atoms of different elements have different sizes and masses.
• Atoms cannot be created, destroyed, or divided into smaller particles.
• Compounds are formed when the atoms of different elements combine with each other in fixed, whole-number ratios.
• Atoms can be combined, separated, or rearranged via chemical reactions.

List two merits of Dalton’s atomic theory.

One of the most important merits of Dalton’s atomic theory is the fact that the theory does not violate several fundamental laws of chemical combination such as the law of definite proportions, the law of multiple proportions, and the law of conservation of mass. Another important merit of Dalton’s atomic theory is that it provided a basis for scientists to differentiate between elements and compounds.

What are the shortcomings of Dalton’s atomic theory?

Some important demerits of Dalton’s atomic theory are listed below.

• The theory did not account for the existence of subatomic particles (it suggested that atoms are indivisible).
• By suggesting that all atoms of an element must have identical masses and sizes, Dalton’s atomic theory did not account for the existence of isotopes. Furthermore, this theory also did not account for the existence of isobars (nuclides of different chemical elements with the same mass number).
• Dalton’s atomic theory failed to explain the dissimilarities in the properties of different allotropes of an element.
• This theory states that elements must combine in simple, whole-number ratios to form compounds. However, this is not necessarily true. Several complex organic compounds do not feature simple ratios of their constituent elements.

Do electrons actually exist?

Most of us realize that the neutron, in an atom of matter, is a negatively charged particle orbiting the nucleus. No two electrons at the same time will occupy the same space. They are part of any molecule, but they may also live on their own, independently.

Which atomic model is used today?

The Bohr paradigm, generally speaking, encapsulates the popular understanding of the atom. In artwork that depicts a central atomic nucleus and oval lines reflecting the electron orbits, this image is also portrayed.

Why can’t you see an atom with the naked eye?

Can atoms be divided or destroyed, what are atoms made of.

Put your understanding of this concept to test by answering a few MCQs. Click ‘Start Quiz’ to begin!

Select the correct answer and click on the “Finish” button Check your score and answers at the end of the quiz

Visit BYJU’S for all Chemistry related queries and study materials

Your result is as below

Request OTP on Voice Call

Leave a Comment Cancel reply

Your Mobile number and Email id will not be published. Required fields are marked *

Post My Comment

Good explanation of questions

It helps to understand the question answer and explain the topic

What is the correction of Dalton’s Atomic theory?

The indivisibility of an atom was proved wrong: an atom can be further subdivided into protons, neutrons and electrons. However, an atom is the smallest particle that takes part in chemical reactions. According to Dalton, the atoms of same element are similar in all respects.

• Share Share

Register with BYJU'S & Download Free PDFs

Register with byju's & watch live videos.

Want to create or adapt books like this? Learn more about how Pressbooks supports open publishing practices.

5.1 Early Atomic Theory: Dalton’s Model of the Atom

Learning objectives.

By the end of this section, you will be able to:

• Summarize the postulates of Dalton’s atomic theory
• Apply Dalton’s atomic theory to explain the laws of definite and multiple proportions

The language used in chemistry is seen and heard in many disciplines, ranging from medicine to engineering to forensics to art. The language of chemistry includes its own vocabulary as well as its own form of shorthand. Chemical symbols are used to represent atoms and elements. Chemical formulas depict molecules as well as the composition of compounds. Chemical equations provide information about the quality and quantity of the changes associated with chemical reactions.

This chapter will lay the foundation for our study of the language of chemistry. The concepts of this foundation include the atomic theory, the composition and mass of an atom, and the variability of the composition of isotopes.

Atomic Theory through the Nineteenth Century

Watch The 2,400-year search for the atom – Theresa Doud (6 mins)

Video source: TED-Ed. (2014, December 8). The 2,400-year search for the atom – Theresa Doud [Video]. YouTube.

The earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers, the scientists of their day. In the fifth century BC, Leucippus and Democritus argued that all matter was composed of small, finite particles that they called atomos , a term derived from the Greek word for “indivisible.” They thought of atoms as moving particles that differed in shape and size, and which could join together. Later, Aristotle and others came to the conclusion that matter consisted of various combinations of the four “elements”—fire, earth, air, and water—and could be infinitely divided. Interestingly, these philosophers thought about atoms and “elements” as philosophical concepts, but apparently never considered performing experiments to test their ideas.

The Aristotelian view of the composition of matter held sway for over two thousand years, until English schoolteacher John Dalton helped to revolutionize chemistry with his hypothesis that the behaviour of matter could be explained using an atomic theory. First published in 1807, many of Dalton’s hypotheses about the microscopic features of matter are still valid in modern atomic theory. Here are the postulates of Dalton’s atomic theory .

• Matter is composed of exceedingly small, indivisible particles called atoms. An atom is the smallest unit of an element that can participate in a chemical change.

• Atoms of one element differ in properties from atoms of all other elements.

• Atoms are neither created nor destroyed during a chemical change, but are instead rearranged and regrouped to yield substances that are different from those present before the change ( Figure 5.1c ).

Dalton’s atomic theory laid the foundation in the development of chemistry. Most of his postulates remain valid; however, some of his conclusions have been revolutionized because further investigations have shown that Atoms are composed of subatomic particles; they are not indivisible Not all the atoms of a specific element have the exact same mass; an element can exist in different forms, called isotopes Atoms, under special conditions, can be decomposed. Source: (Hein & Arena, 2014, p. 83)

Despite these flaws, Dalton’s atomic theory provides a microscopic explanation of the many macroscopic properties of matter that you’ve learned about. For example, if an element such as copper consists of only one kind of atom, then it cannot be broken down into simpler substances, that is, into substances composed of fewer types of atoms. And if atoms are neither created nor destroyed during a chemical change, then the total mass of matter present when matter changes from one type to another will remain constant (the law of conservation of matter).

Example 5.1a

Testing dalton’s atomic theory.

In the following drawing, the green, larger spheres represent atoms of a certain element. The blue, smaller spheres represent atoms of another element. If the spheres touch, they are part of a single unit of a compound. Does the following chemical change represented by these symbols violate any of the ideas of Dalton’s atomic theory? If so, which one?

The starting materials consist of two green, larger spheres and two blue, smaller spheres. The products consist of only one green sphere and one blue sphere. This violates Dalton’s postulate that atoms are neither created nor destroyed during a chemical change, but are merely redistributed. (In this case, atoms appear to have been destroyed.)

Exercise 5.1a

Check Your Answer

Exercise 5.1b.

Choose the answer that best answers the questions for each of the multiple choice questions.

• Subatomic particles
• Law of constant proportions
• Law of conservation of mass
• Law of multiple proportions
• Law of electric force
• Temperature, volume
• Force, pressure
• Weight, volume
• Physical and chemical properties, mass
• Atoms that combine to form new molecules do so in simple, whole number ratios
• A chemical reaction is a rearrangement of atoms. No atoms are created or destroyed.
• All elements are composed of small particles called atoms.
• Atoms of a given element are always identical.
• Atoms are always on motion
• Matter is composed of exceedingly small, indivisible particles.
• Elements consist of only one type of identical atom, which has the same mass for all atoms.
• Theories stated in a) and b) were both proven to be false
• Neither theory stated in a) nor b) were proven to be false

Source: “Exercise 5.1b” by Jackie MacDonald, licensed under CC BY-NC-SA 4.0

Dalton knew of the experiments of French chemist Joseph Proust, who demonstrated that all samples of a pure compound contain the same elements in the same proportion by mass . This statement is known as the law of definite proportions or the law of constant composition . The suggestion that the numbers of atoms of the elements in a given compound always exist in the same ratio is consistent with these observations. For example, when different samples of isooctane (a component of gasoline and one of the standards used in the octane rating system) are analyzed, they are found to have a carbon-to-hydrogen mass ratio of 5.33:1, as shown in Table 5.1a.

It is worth noting that although all samples of a particular compound have the same mass ratio, the converse is not true in general. That is, samples that have the same mass ratio are not necessarily the same substance. For example, there are many compounds other than isooctane that also have a carbon-to-hydrogen mass ratio of 5.33 : 1.00.

Watch Thinking Reeds Chemistry – The Law of Definite Proportions (4 mins 13 sec)

Video Source: Thinking Reeds Chemistry (2016, September 6). Thinking Reeds Chemistry – The Law of Definite Proportions [Video]. YouTube.

Dalton also used data from Proust, as well as results from his own experiments, to formulate another interesting law. The law of multiple proportions states that when two elements react to form more than one compound, a fixed mass of one element will react with masses of the other element in a ratio of small, whole numbers . For example, copper and chlorine can form a green, crystalline solid with a mass ratio of 0.558 g chlorine to 1 g copper, as well as a brown crystalline solid with a mass ratio of 1.116 g chlorine to 1 g copper. These ratios by themselves may not seem particularly interesting or informative; however, if we take a ratio of these ratios, we obtain a useful and possibly surprising result: a small, whole-number ratio.

This 2-to-1 ratio means that the brown compound has twice the amount of chlorine per amount of copper as the green compound. When referencing Figure 5.1d, the above can be explained by atomic theory if the copper-to-chlorine ratio in the brown compound (Figure 5.1d (b)) is 1 copper atom to 2 chlorine atoms, and the ratio in the green compound (Figure 5.1d (a)) is 1 copper atom to 1 chlorine atom. The ratio of chlorine atoms in compound B compared to compound A (and thus the ratio of their masses) is therefore 2 to 1 (Figure 5.1d).

Watch Multiple Proportions (1 min 11 sec)

Video Source: kwesiamoa. (2010, January 20). Multiple Proportions [Video]. YouTube.

Example 5.1b

Laws of definite and multiple proportions.

A sample of compound A (a clear, colourless gas) is analyzed and found to contain 4.27 g carbon and 5.69 g oxygen. A sample of compound B (also a clear, colourless gas) is analyzed and found to contain 5.19 g carbon and 13.84 g oxygen. Are these data an example of the law of definite proportions, the law of multiple proportions, or neither? What do these data tell you about substances A and B?

In compound A, the mass ratio of carbon to oxygen is:

In compound B, the mass ratio of carbon to oxygen is:

The ratio of these ratios is:

The ratio of oxygen atoms of compound A to compound B (and thus the ratio of their masses) is 1 to 2. This supports the law of multiple proportions. This means that A and B are different compounds, with A having one-half as much oxygen per amount of carbon as compound B. A possible pair of compounds that would fit this relationship would be A = CO and B = CO 2 .

Exercise 5.1c

A sample of compound X (a clear, colourless, combustible liquid with a noticeable odour) is analyzed and found to contain 14.13 g carbon and 2.96 g hydrogen. A sample of compound Y (a clear, colourless, combustible liquid with a noticeable odour that is slightly different from X’s odour) is analyzed and found to contain 19.91 g carbon and 3.34 g hydrogen. Are these data an example of the law of definite proportions, the law of multiple proportions, or neither? What do these data tell you about substances X and Y?

Links to Interactive Learning Tools

Explore the Timeline of Atomic Discovery  from eCampusOntario H5P Studio .

Attribution & References

Except where otherwise noted, this page is adapted by Jackie MacDonald from “ 2.1 Early ideas in atomic theory ” In Chemistry 2e (Open Stax) by Paul Flowers, Klaus Theopold, Richard Langley & William R. Robinson is licensed under CC BY 4.0. Access for free at Chemistry 2e (Open Stax).

Hein, M., & Arena, S. (2014). Foundations of College Chemistry (14th edition) . Wiley & Sons.

• The starting materials consist of four green, larger spheres and two blue, smaller spheres. The products consist of four green, larger spheres and two blue, smaller spheres. This does not violate any of Dalton’s postulates: Atoms are neither created nor destroyed, but are redistributed in small, whole-number ratios. ↵
• 1c; 2b; 3d; 4e; 5c ↵
• In compound X, the mass ratio of carbon to hydrogen is [latex]\frac{14.13 \text{g C}}{2.96 \text{g H}}[/latex]. In compound Y, the mass ratio of carbon to oxygen is [latex]\frac{19.91 \text{g C}}{3.34 \text{g H}}[/latex]. The ratio of these ratios is [latex]\frac{\frac{14.13 \text{g C}}{2.96 \text{g H}}}{\frac{19.91 \text{g C}}{3.34 \text{g H}}} = \frac{4.77 \text{g C/g H}}{5.96 \text{g C/g H}} = 0.800 = \frac{4}{5}[/latex]. This small, whole-number ratio supports the law of multiple proportions. This means that X and Y are different compounds. ↵

set of postulates that established the fundamental properties of atoms

smallest particle of an element that can enter into a chemical combination

substance that is composed of a single type of atom; a substance that cannot be decomposed by a chemical change

(also, law of constant composition) all samples of a pure compound contain the same elements in the same proportions by mass

(also, law of definite proportions) all samples of a pure compound contain the same elements in the same proportions by mass

when two elements react to form more than one compound, a fixed mass of one element will react with masses of the other element in a ratio of small whole numbers

Enhanced Introductory College Chemistry Copyright © 2023 by Gregory Anderson; Caryn Fahey; Jackie MacDonald; Adrienne Richards; Samantha Sullivan Sauer; J.R. van Haarlem; and David Wegman is licensed under a Creative Commons Attribution 4.0 International License , except where otherwise noted.

Share This Book

• Privacy Policy

Read Chemistry

Dalton’s atomic theory: definition, statement, and postulates.

General Chemistry

– In this subject, we will discuss the Dalton’s Atomic Theory: Definition, Statement, and Postulates

Atomic Theory

– In this lesson, we will explain atomic Theory especially, Dalton’s atomic theory.

– In the fifth-century b.c., the Greek philosopher Democritus expressed the belief that all matter consists of very small, indivisible particles. he named atomos (meaning uncuttable or indivisible).

– Although Democritus’ idea was not accepted by many of his contemporaries (notably Plato and Aristotle), somehow it endured.

– Experimental evidence from early scientific investigations provided support for the notion of “atomism” and gradually gave rise to the modern definitions of elements and compounds.

– It was in 1808 that an English scientist and schoolteacher, John Dalton, formulated a precise definition of the indivisible building blocks of matter that we call atoms.

Dalton’s atomic theory

– Dalton’s work marked the beginning of the modern era of chemistry.

– The hypotheses about the nature of matter on which Dalton’s atomic theory is based can be summarized as:

(1)   Elements are composed of extremely small particles, called atoms.

(2)   All atoms of a given element are identical, having the same size, mass, and chemical properties.

– The atoms of one element are different from the atoms of all other elements.

(3) Compounds are composed of atoms of more than one element.

– In any compound, the ratio of the number of atoms of any two of the elements present is either an integer or a simple fraction.

(4)  A chemical reaction involves only the separation, combination, or rearrangement of atoms.

– it does not result in their creation or destruction.

– In this figure:

(a)  According to Dalton’s atomic theory, atoms of the same element are identical, but atoms of one element are different from atoms of other elements.

(b)  Compounds formed from atoms of elements X and Y.

– In this case, the ratio of the atoms of element X to the atoms of element Y is 2:1.

– The Figure above is a schematic representation of hypotheses 2 and 3.

– Dalton’s concept of an atom was far more detailed and specific than Democritus’.

Dalton’s second hypothesis

– The second hypothesis states that atoms of one element are different from atoms of all other elements.

– Dalton did not attempt to describe the structure or composition of atoms. he had no idea what an atom was really like.

– However, he did realize that the different properties shown by elements such as hydrogen and oxygen can be explained by assuming that hydrogen atoms are not the same as oxygen atoms.

Dalton’s third hypothesis

– The third hypothesis suggests that to form a certain compound, we need not only atoms of the right kinds of elements. But the specific numbers of these atoms as well.

– This idea is an extension of a law published in 1799 by Joseph Proust, a French chemist.

– Proust’s law of definite proportions states that different samples of the same compound always contain its constituent elements in the same proportion by mass.

– Thus, if we were to analyze samples of carbon dioxide gas obtained from different sources, we would find in each sample the same ratio by mass of carbon to oxygen.

– It stands to reason, then, that if the ratio of the masses of different elements in a given compound is fixed, the ratio of the atoms of these elements in the compound also must be constant.

– Dalton’s third hypothesis also supports another important law, the  law of multiple proportions .

– According to this law, if two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers.

– Dalton’s theory explains the law of multiple proportions quite simply: The compounds differ in the number of atoms of each kind that combine.

Example of the law of multiple proportions

– For example, carbon forms two stable compounds with oxygen, namely, carbon monoxide and carbon dioxide.

– Modern measurement techniques indicate that one atom of carbon combines with one atom of oxygen in carbon monoxide and that one atom of carbon combines with two oxygen atoms in carbon dioxide.

– Thus, the ratio of oxygen in carbon monoxide to oxygen in carbon dioxide is 1:2.

– This result is consistent with the law of multiple proportions because the mass of an element in a compound is proportional to the number of atoms of the element present.

Dalton’s fourth hypothesis

– Dalton’s fourth hypothesis is another way of stating the law of conservation of mass, which is that matter can be neither created nor destroyed.

– Because matter is made of atoms that are unchanged in a chemical reaction, it follows that mass must be conserved as well.

– Dalton’s brilliant insight into the nature of matter was the main stimulus for the rapid progress of chemistry during the nineteenth century.

Conclusion of Dalton’s atomic theory

(1) we can define an atom as the basic unit of an element that can enter into a chemical combination. 

(2) Dalton imagined an atom that was both extremely small and indivisible.

– However, a series of investigations that began in the 1850s and extended into the twentieth century demonstrated that atoms possess internal structure. 

(3) Atoms are made up of even smaller particles, which are called subatomic particles.

– This research led to the discovery of three such particles— electrons, protons, and neutrons .

Reference:  General Chemistry: The Essential Concepts / Raymond Chang, Jason Overby. ( sixth edition).

Related Articles

Molecular Orbitals for Heteronuclear Diatomic Molecules

May 20, 2019

Molecular Orbitals for Homonuclear Diatomic Molecules

MCQ on Chemical Bonding – Orbital Theory

Molecular Orbital Theory

MCQ on Chapter: structure of atom – wave mechanical approach

May 19, 2019

Electronegativity and Electron Affinity

Leave a reply cancel reply.

Your email address will not be published. Required fields are marked *

Save my name, email, and website in this browser for the next time I comment.

Dalton's Atomic Theory | General Chemistry 1

What are the 5 points of dalton's atomic theory.

Dalton's Atomic Theory was formulated by John Dalton in 1808, and it remains a fundamental tenet of chemistry to this day. The five main points are:

• Matter is made up of atoms, small and indivisible particles.
• All atoms of the same element are identical and have the same mass.
• Atoms of different elements vary in size, mass, and chemical behavior.
• Chemical compounds are made up of at least 2 atoms of different elements. The resulting particle is called a molecule.
• In a chemical reaction, atoms are rearranged, separated, or recombined to form new compounds but no atom is created or destroyed.

While Dalton was considered a pioneer of modern chemistry, some of his theories were later proven inaccurate. Despite this, many of his ideas became the foundation for discoveries in chemistry in the coming decades.

Which Part of Dalton's Theory Was Wrong?

The indivisibility of an atom is no longer accepted, as it was discovered that atoms can be further divided into protons, neutrons, and electrons. 

• Dalton also proposed that atoms of different elements combine in simple whole-number ratios such as 2:3 or 3:5 for Hydrogen and Oxygen respectively. In contrast, complex organic compounds like sugar (C 12 H 22 O 11 ) defy this theory.
• Dalton's theories were also wrong when it came to, for example, calcium and argon. In the case of these two elements, they share a 40 amu atomic mass. This means Dalton's theory of atoms of different elements being different in all respects was incorrect in certain cases.
• Atoms of the same element typically will have similar properties as Dalton proposed. However, there are a few exceptions. These are known as isotopes. Isotopes are atoms with different masses, which goes against Dalton's original theory. For example, chlorine has two isotopes: one at mass number 35 and another at 37.
• Finally, allotropes are a poorly explained phenomenon that cannot be accounted for by Dalton's theory. Charcoal, graphite, and diamond have many differences in properties from one another despite having similar structures which suggest there may exist an underlying explanation as to why this happens.

Who Proved Dalton's Theory Wrong?

In a world-shattering discovery, English physicist J.J Thomson disproved the idea that atoms are indivisible.  Thomson discovered that an atom consists of a heavy and dense core, known as the nucleus, which is orbited by much lighter particles: protons and electrons.

How Did Thomson Prove This?

Thomson entered a modified Crookes tube into an electrical field. He noticed that the tube gave off a green glow as cathode rays hit the opposite electrode. Thomson concluded that the glowing particles were not simply light, but had mass because they were deflected in opposite directions by an electrical field.

Thomson realized that cathode rays could be made up of only one kind of particle if he was able to deflect them using an electric field without changing their properties, which was possible with a magnet. When the rays were deflected by a magnet, Thomson concluded that they must be made up of negatively charged particles since only particles with negative charges would curve in the opposite direction to an electric field.

Thomson also found that these particles could be both absorbed and released by the element neon, which meant they were not stable atoms. This led Thomson to conclude that cathode rays were streams of particles, each with a negative charge and mass.

Thomson's discoveries led him to propose the first complete model of the atom: an atom has a positively charged nucleus surrounded by negatively charged electrons in orbit. 

• Sign in with Google
• Sign in with Facebook

Distillations magazine

John dalton and the scientific method.

Dalton proposed atomic theory in 1808; an additional century passed before the theory was universally accepted by scientists.

Many consider 2008 the 200th anniversary of atomic theory, John Dalton’s momentous theory of the nature of matter. Dalton (1766–1844) proposed that all matter in the universe is made of indestructible, unchangeable atoms—each type characterized by a constant mass—that undergo chemical reactions by joining with and separating from each other. But anniversaries can be deceptive. It was 1808 when Dalton published the first volume of New Systems of Chemical Philosophy , which presented his atomic theory in full, but his ideas were in fact already known, as he had been talking and writing about them for at least five years. Yet, an additional century would pass before atomic theory became universally accepted. 

The theory certainly had its early fans, including Swedish chemist Jöns Jakob Berzelius (1779–1848). There was hard evidence in its favor; conceiving of atoms in this way explained the stoichiometry of reactions, which posited that combined elements retained their proportions before, during, and after reacting with each other. However, not everyone found this fact compelling. Humphry Davy (British, 1748–1829) and Claude-Louis Berthollet (French, 1748–1822) were not convinced. Because atoms could not be seen, Dalton could not base his theory on direct observation, and this was a major stumbling block for many scientists. 

Nevertheless atomic theory was useful, whether proven or not. It was easier to express stoichiometric proportions in terms of atoms than in terms of absolute mass. It is simpler to say that 1 atom of hydrogen joins with 1 atom of chlorine to form 1 molecule of hydrogen chloride than it is to say that 1 gram of hydrogen reacts with 35.45 grams of chlorine to make 36.45 grams of hydrogen chloride. Many chemists found themselves using atomic theory, even if they held their noses all the while. 

Acceptance grew slowly over the next hundred years as the concept of the atom became useful for explaining a variety of things from molecular structure in organic chemistry to the spacing and movement of molecules in gas physics. By 1905 there were still some holdouts, including Marcellin Berthelot and the founding father of physical chemistry, Wilhelm Ostwald, but most chemists had accepted the existence of atoms. That year a young Albert Einstein penned a paper that doesn’t receive nearly as much attention as his work on the photoelectric effect and his special theory of relativity. This work used the concept of the atom to explore the phenomenon of Brownian motion. 

When minute particles are suspended in a liquid, they move in a seemingly random, ever-changing course, each one only slowly moving in any direction. Earlier scientists had proposed that the particles moved because the liquid molecules were constantly in motion and collided with the suspended particles, jostling them in an erratic manner. Einstein took this idea further, building on the observation of Jacobus Henricus van’t Hoff that solute molecules move in the same manner as gas molecules and their behavior can be described using the gas laws. In his 1905 paper Einstein treated suspended particles as if they were giant molecules and went on to predict how they should behave according to the gas laws. For example, he stated that the average speed of the suspended particles should reflect the average kinetic energy of the moving molecules of the liquid in which the particles were suspended. He also predicted that, in a vertical cylinder of an aqueous suspension, gravitational pull would cause a greater density of particles toward the bottom of the cylinder and a lower density near the top, just as the earth’s atmosphere becomes thinner at higher altitudes. 

Three years later, in 1908, Dalton’s New System of Chemical Philosophy turned 100 years old. Berthelot had died the previous year, rejecting atoms until the end. Ostwald still did not accept the existence of atoms. French scientist Jean Perrin took up Einstein’s challenge and began studying Brownian motion in great detail. Perrin carried out incredibly meticulous observations, plotting the paths of protein particles in aqueous suspensions. He studied their variations in distribution as a function of the tiniest variations in vertical height. He was able to show that their behavior matched Einstein’s predictions for particles that are being constantly rammed by unseen molecules. Once and for all the particulate nature of matter had been demonstrated in an unequivocal manner. A final crowning validation came in 1926 when Perrin received the Nobel Prize in Physics for his work. 

Later scientists would use atomic force microscopy and scanning tunneling microscopy to give us even clearer observations of the particulate nature of matter. Today atomic theory is covered in the first chapters of most general chemistry textbooks. Most of us have grown accustomed to seeing it in this exalted place, and it can be easy to forget that it has not been there all along since 1808.

Mark Michalovic was consultant for educational services at the Science History Institute.

More from our magazine

Sylvia Earle and the Call of the Deep

Adventure and tangled interests under the sea.

Matchmaking in Colonial India

An inconspicuous technology sparks revolution on the subcontinent.

The Eclipse That Killed a King (and May Have Saved a Kingdom)

How the scientific prowess of King Mongkut of Siam helped stave off European incursion.

Copy the above HTML to republish this content. We have formatted the material to follow our guidelines, which include our credit requirements. Please review our full list of guidelines for more information. By republishing this content, you agree to our republication requirements.

• Subscriber Services
• For Authors
• Publications
• Archaeology
• Art & Architecture
• Bilingual dictionaries
• Classical studies
• Encyclopedias
• English Dictionaries and Thesauri
• Language reference
• Linguistics
• Media studies
• Medicine and health
• Names studies
• Performing arts
• Science and technology
• Social sciences
• Society and culture
• Overview Pages
• Subject Reference
• English Dictionaries
• Bilingual Dictionaries

Recently viewed (0)

• Save Search
• Share This Facebook LinkedIn Twitter

Related Content

More like this.

Show all results sharing these subjects:

Dalton's atomic theory

Quick reference.

A theory of chemical combination, first stated by John Dalton in 1803. It involves the following postulates: (1) Elements consist of indivisible small particles (atoms). (2) All atoms of the same element are identical; different elements have different types of atom. (3) Atoms can neither be created nor destroyed. (4) ‘Compound elements’ (i.e. compounds) are formed when atoms of different elements join in simple ratios to form ‘compound atoms’ (i.e. molecules).Dalton also proposed symbols for atoms of different elements (later replaced by the present notation using letters).

(1) Elements consist of indivisible small particles (atoms).

(2) All atoms of the same element are identical; different elements have different types of atom.

(3) Atoms can neither be created nor destroyed.

(4) ‘Compound elements’ (i.e. compounds) are formed when atoms of different elements join in simple ratios to form ‘compound atoms’ (i.e. molecules).

From:   Dalton's atomic theory   in  A Dictionary of Physics »

Subjects: Science and technology — Chemistry

Related content in Oxford Reference

Reference entries.

View all related items in Oxford Reference »

Search for: 'Dalton's atomic theory' in Oxford Reference »

• Oxford University Press

PRINTED FROM OXFORD REFERENCE (www.oxfordreference.com). (c) Copyright Oxford University Press, 2023. All Rights Reserved. Under the terms of the licence agreement, an individual user may print out a PDF of a single entry from a reference work in OR for personal use (for details see Privacy Policy and Legal Notice ).

date: 28 April 2024

• Cookie Policy
• Privacy Policy
• Legal Notice
• Accessibility
• [66.249.64.20|185.147.128.134]
• 185.147.128.134

Character limit 500 /500

John Dalton's Atomic Theory

• Famous Chemists
• Chemical Laws
• Periodic Table
• Projects & Experiments
• Scientific Method
• Biochemistry
• Physical Chemistry
• Medical Chemistry
• Chemistry In Everyday Life
• Activities for Kids
• Abbreviations & Acronyms
• Weather & Climate
• Ph.D., Biomedical Sciences, University of Tennessee at Knoxville
• B.A., Physics and Mathematics, Hastings College

You may take it for granted that matter is made up of atoms , but what we consider common knowledge was unknown until relatively recently in human history. Most science historians credit John Dalton , a British physicist, chemist, and meteorologist, with the development of modern atomic theory.

Early Theories 

While the ancient Greeks believed atoms made matter, they disagreed on what atoms were. Democritus recorded that Leucippus believed atoms to be small, indestructible bodies that could combine to change properties of matter. Aristotle believed elements each had their own special "essence," but he did not think the properties extended down to tiny, invisible particles. No one really questioned Aristotle's theory, since tools did not exist to examine matter in detail.

Along Comes Dalton

So, it wasn't until the 19th century that scientists conducted experiments on the nature of matter. Dalton's experiments focused on gases -- their properties, what happened when they were combined, and the similarities and differences between different types of gases. What he learned led him to propose several laws, which are known collectively as Dalton's Atomic Theory or Dalton's Laws:

• Atoms are small, chemically indestructible particles of matter. Elements consist of atoms.
• Atoms of an element share common properties.
• Atoms of different elements have different properties and different atomic weights.
• Atoms that interact with each other obey the Law of Conservation of Mass . Essentially, this law states the number and kinds of atoms that react are equal to the number and kinds of atoms in the products of a chemical reaction.
• Atoms that combine with each other obey the Law of Multiple Proportions . In other words, when elements combine, the ratio in which the atoms combine can be expressed as a ratio of whole numbers.

Dalton is also known for proposing gas laws ( Dalton's Law of Partial Pressures ) and explaining color blindness. Not all of his scientific experiments could be called successful. For example, some believe the stroke he suffered might have resulted from research using himself as a subject, in which he poked himself in the ear with a sharp stick to “investigate the humours that move inside of my cranium.”

• Grossman, M. I. (2014). "John Dalton and the London atomists: William and Bryan Higgins, William Austin, and new Daltonian doubts about the origin of the atomic theory." Notes and Records . 68 (4): 339–356. doi: 10.1098/rsnr.2014.0025
• Levere, Trevor (2001). Transforming Matter: A History of Chemistry from Alchemy to the Buckyball . Baltimore, Maryland: The Johns Hopkins University Press. pp. 84–86. ISBN 978-0-8018-6610-4.
• Rocke, Alan J. (2005). "In Search of El Dorado: John Dalton and the Origins of the Atomic Theory." Social Research. 72 (1): 125–158. JSTOR 40972005
• A Brief History of Atomic Theory
• Biography of John Dalton, the 'Father of Chemistry'
• Law of Multiple Proportions Example Problem
• Law of Constant Composition in Chemistry
• Law of Definite Proportions Definition
• Basic Model of the Atom and Atomic Theory
• Atoms and Atomic Theory - Study Guide
• The Major Laws of Chemistry
• Gases - General Properties of Gases
• J.J. Thomson Atomic Theory and Biography
• What Is Dalton's Law of Partial Pressures?
• Biography of Amedeo Avogadro, Influential Italian Scientist
• Gases Study Guide
• A to Z Chemistry Dictionary
• 10 Famous Meteorologists
• Overview of High School Chemistry Topics

2.1 Early Ideas in Atomic Theory

Learning objectives.

• State the postulates of Dalton’s atomic theory
• Use postulates of Dalton’s atomic theory to explain the laws of definite and multiple proportions

The language used in chemistry is seen and heard in many disciplines, ranging from medicine to engineering to forensics to art. The language of chemistry includes its own vocabulary as well as its own form of shorthand. Chemical symbols are used to represent atoms and elements. Chemical formulas depict molecules as well as the composition of compounds. Chemical equations provide information about the quality and quantity of the changes associated with chemical reactions.

This chapter will lay the foundation for our study of the language of chemistry. The concepts of this foundation include the atomic theory, the composition and mass of an atom, the variability of the composition of isotopes, ion formation, chemical bonds in ionic and covalent compounds, the types of chemical reactions, and the naming of compounds. We will also introduce one of the most powerful tools for organizing chemical knowledge: the periodic table.

Atomic Theory through the Nineteenth Century

The earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers, the scientists of their day. In the fifth century BC, Leucippus and Democritus argued that all matter was composed of small, finite particles that they called atomos , a term derived from the Greek word for “indivisible.” They thought of atoms as moving particles that differed in shape and size, and which could join together. Later, Aristotle and others came to the conclusion that matter consisted of various combinations of the four “elements”—fire, earth, air, and water—and could be infinitely divided. Interestingly, these philosophers thought about atoms and “elements” as philosophical concepts, but apparently never considered performing experiments to test their ideas.

The Aristotelian view of the composition of matter held sway for over two thousand years, until English schoolteacher John Dalton helped to revolutionize chemistry with his hypothesis that the behavior of matter could be explained using an atomic theory. First published in 1807, many of Dalton’s hypotheses about the microscopic features of matter are still valid in modern atomic theory. Here are the postulates of Dalton’s atomic theory .

• Matter is composed of exceedingly small particles called atoms. An atom is the smallest unit of an element that can participate in a chemical change.
• Atoms of one element differ in properties from atoms of all other elements.

Dalton’s atomic theory provides a microscopic explanation of the many macroscopic properties of matter that you’ve learned about. For example, if an element such as copper consists of only one kind of atom, then it cannot be broken down into simpler substances, that is, into substances composed of fewer types of atoms. And if atoms are neither created nor destroyed during a chemical change, then the total mass of matter present when matter changes from one type to another will remain constant (the law of conservation of matter).

Example 2.1

Testing dalton’s atomic theory, check your learning.

The starting materials consist of four green spheres and two purple spheres. The products consist of four green spheres and two purple spheres. This does not violate any of Dalton’s postulates: Atoms are neither created nor destroyed, but are redistributed in small, whole-number ratios.

Dalton knew of the experiments of French chemist Joseph Proust, who demonstrated that all samples of a pure compound contain the same elements in the same proportion by mass . This statement is known as the law of definite proportions or the law of constant composition . The suggestion that the numbers of atoms of the elements in a given compound always exist in the same ratio is consistent with these observations. For example, when different samples of isooctane (a component of gasoline and one of the standards used in the octane rating system) are analyzed, they are found to have a carbon-to-hydrogen mass ratio of 5.33:1, as shown in Table 2.1 .

It is worth noting that although all samples of a particular compound have the same mass ratio, the converse is not true in general. That is, samples that have the same mass ratio are not necessarily the same substance. For example, there are many compounds other than isooctane that also have a carbon-to-hydrogen mass ratio of 5.33:1.00.

Dalton also used data from Proust, as well as results from his own experiments, to formulate another interesting law. The law of multiple proportions states that when two elements react to form more than one compound, a fixed mass of one element will react with masses of the other element in a ratio of small, whole numbers . For example, copper and chlorine can form a green, crystalline solid with a mass ratio of 0.558 g chlorine to 1 g copper, as well as a brown crystalline solid with a mass ratio of 1.116 g chlorine to 1 g copper. These ratios by themselves may not seem particularly interesting or informative; however, if we take a ratio of these ratios, we obtain a useful and possibly surprising result: a small, whole-number ratio.

This 2-to-1 ratio means that the brown compound has twice the amount of chlorine per amount of copper as the green compound.

This can be explained by atomic theory if the copper-to-chlorine ratio in the brown compound is 1 copper atom to 2 chlorine atoms, and the ratio in the green compound is 1 copper atom to 1 chlorine atom. The ratio of chlorine atoms (and thus the ratio of their masses) is therefore 2 to 1 ( Figure 2.5 ).

Example 2.2

Laws of definite and multiple proportions.

In compound B, the mass ratio of carbon to oxygen is:

The ratio of these ratios is:

This supports the law of multiple proportions. This means that A and B are different compounds, with A having one-half as much oxygen per amount of carbon (or twice as much carbon per amount of oxygen) as B. A possible pair of compounds that would fit this relationship would be A = CO and B = CO 2 .

In compound X, the mass ratio of carbon to hydrogen is 14.13 g C 2.96 g H . 14.13 g C 2.96 g H . In compound Y, the mass ratio of carbon to oxygen is 19.91 g C 3.34 g H . 19.91 g C 3.34 g H . The ratio of these ratios is 14.13 g C 2.96 g H 19.91 g C 3.34 g H = 4.77 g C/g H 5.96 g C/g H = 0.800 = 4 5 . 14.13 g C 2.96 g H 19.91 g C 3.34 g H = 4.77 g C/g H 5.96 g C/g H = 0.800 = 4 5 . This small, whole-number ratio supports the law of multiple proportions. This means that X and Y are different compounds.

As an Amazon Associate we earn from qualifying purchases.

This book may not be used in the training of large language models or otherwise be ingested into large language models or generative AI offerings without OpenStax's permission.

Want to cite, share, or modify this book? This book uses the Creative Commons Attribution License and you must attribute OpenStax.

Access for free at https://openstax.org/books/chemistry/pages/1-introduction

• Authors: Paul Flowers, William R. Robinson, PhD, Richard Langley, Klaus Theopold
• Publisher/website: OpenStax
• Book title: Chemistry
• Publication date: Mar 11, 2015
• Location: Houston, Texas
• Book URL: https://openstax.org/books/chemistry/pages/1-introduction
• Section URL: https://openstax.org/books/chemistry/pages/2-1-early-ideas-in-atomic-theory

© Feb 15, 2022 OpenStax. Textbook content produced by OpenStax is licensed under a Creative Commons Attribution License . The OpenStax name, OpenStax logo, OpenStax book covers, OpenStax CNX name, and OpenStax CNX logo are not subject to the Creative Commons license and may not be reproduced without the prior and express written consent of Rice University.

• school Campus Bookshelves
• menu_book Bookshelves
• perm_media Learning Objects
• login Login
• how_to_reg Request Instructor Account
• hub Instructor Commons
• Download Page (PDF)
• Download Full Book (PDF)
• Periodic Table
• Physics Constants
• Scientific Calculator
• Reference & Cite
• Tools expand_more
• Readability

selected template will load here

This action is not available.

2.1: Historical Development of Atomic Theory

• Last updated
• Save as PDF
• Page ID 330284

Notes about this page

The page below is a brief overview on the history of atomic theory. It contains lots of video media .

Alternatives pages on the history of atomic theory are:

• A more detailed overview that includes practice problems is here (click) .
• Another brief overview is here (click) .

Skills to Develop

By the end of this section, you will be able to: 

• State the postulates of Dalton’s atomic theory
• Use postulates of Dalton’s atomic theory to explain the laws of definite and multiple proportions
• Outline milestones in the development of modern atomic theory
• Summarize and interpret the results of the experiments of Thomson, Millikan, and Rutherford

A Video Introduction to Atomic Theory through the Nineteenth Century From Crash Course Chemistry

Video \(\PageIndex{1}\): Lavoisier's discovery of The Law of Conservation of Matter led to the Laws of Definite and Multiple Proportions and eventually Dalton's Atomic Theory.

Atomic Theory through the Nineteenth Century

The earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers, the scientists of their day. In the fifth century BC, Leucippus and Democritus argued that all matter was composed of small, finite particles that they called atomos , a term derived from the Greek word for “indivisible.” They thought of atoms as moving particles that differed in shape and size, and which could join together. Later, Aristotle and others came to the conclusion that matter consisted of various combinations of the four “elements”—fire, earth, air, and water—and could be infinitely divided. Interestingly, these philosophers thought about atoms and “elements” as philosophical concepts, but apparently never considered performing experiments to test their ideas.

The Aristotelian view of the composition of matter held sway for over two thousand years, until English schoolteacher John Dalton helped to revolutionize chemistry with his hypothesis that the behavior of matter could be explained using an atomic theory. First published in 1807, many of Dalton’s hypotheses about the microscopic features of matter are still valid in modern atomic theory. Here are the postulates of Dalton’s atomic theory .

• Matter is composed of exceedingly small particles called atoms. An atom is the smallest unit of an element that can participate in a chemical change.
• An element consists of only one type of atom, which has a mass that is characteristic of the element and is the same for all atoms of that element (Figure \(\PageIndex{1}\)). A macroscopic sample of an element contains an incredibly large number of atoms, all of which have identical chemical properties.
• Atoms of one element differ in properties from atoms of all other elements.
• A compound consists of atoms of two or more elements combined in a small, whole-number ratio. In a given compound, the numbers of atoms of each of its elements are always present in the same ratio (Figure \(\PageIndex{2}\)).
• Atoms are neither created nor destroyed during a chemical change, but are instead rearranged to yield substances that are different from those present before the change (Figure \(\PageIndex{3}\)).

Figure \(\PageIndex{1}\): A pre-1982 copper penny (left) contains approximately 3 \(\times\) 10 22 copper atoms (several dozen are represented as brown spheres at the right), each of which has the same chemical properties. (credit: modification of work by “slgckgc”/Flickr)

Figure \(\PageIndex{2}\): Copper(II) oxide, a powdery, black compound, results from the combination of two types of atoms—copper (brown spheres) and oxygen (red spheres)—in a 1:1 ratio. (credit: modification of work by “Chemicalinterest”/Wikimedia Commons)

Figure \(\PageIndex{3}\) : When the elements copper (a shiny, red-brown solid, shown here as brown spheres) and oxygen (a clear and colorless gas, shown here as red spheres) react, their atoms rearrange to form a compound containing copper and oxygen (a powdery, black solid). (credit copper: modification of work by http://images-of-elements.com/copper.php ).

Dalton’s atomic theory provides a microscopic explanation of the many macroscopic properties of matter that you’ve learned about. For example, if an element such as copper consists of only one kind of atom, then it cannot be broken down into simpler substances, that is, into substances composed of fewer types of atoms. And if atoms are neither created nor destroyed during a chemical change, then the total mass of matter present when matter changes from one type to another will remain constant (the law of conservation of matter (or mass)).

Want to learn more about the Law of Conservation of Mass?

Video \(\PageIndex{2}\): "We are made of star stuff" - Carl Sagan .

Example \(\PageIndex{1}\): Test ing Dalton’s Atomic Theory

In the following drawing, the green spheres represent atoms of a certain element. The purple spheres represent atoms of another element. If the spheres touch, they are part of a single unit of a compound. Does the following chemical change represented by these symbols violate any of the ideas of Dalton’s atomic theory? If so, which one?

The starting materials consist of two green spheres and two purple spheres. The products consist of only one green sphere and one purple sphere. This violates Dalton’s postulate that atoms are neither created nor destroyed during a chemical change, but are merely redistributed. (In this case, atoms appear to have been destroyed.)

Exercise \(\PageIndex{1}\)

In the following drawing, the green spheres represent atoms of a certain element. The purple spheres represent atoms of another element. If the spheres touch, they are part of a single unit of a compound. Does the following chemical change represented by these symbols violate any of the ideas of Dalton’s atomic theory? If so, which one

The starting materials consist of four green spheres and two purple spheres. The products consist of four green spheres and two purple spheres. This does not violate any of Dalton’s postulates: Atoms are neither created nor destroyed, but are redistributed in small, whole-number ratios.

Dalton knew of the experiments of French chemist Joseph Proust, who demonstrated that all samples of a pure compound contain the same elements in the same proportion by mass . This statement is known as the law of definite proportions or the law of constant composition . The suggestion that the numbers of atoms of the elements in a given compound always exist in the same ratio is consistent with these observations. For example, when different samples of isooctane (a component of gasoline and one of the standards used in the octane rating system) are analyzed, they are found to have a carbon-to-hydrogen mass ratio of 5.33:1, as shown in Table \(\PageIndex{1}\).

It is worth noting that although all samples of a particular compound have the same mass ratio, the converse is not true in general. That is, samples that have the same mass ratio are not necessarily the same substance. For example, there are many compounds other than isooctane that also have a carbon-to-hydrogen mass ratio of 5.33:1.00.

Dalton also used data from Proust, as well as results from his own experiments, to formulate another interesting law. The law of multiple proportions states that when two elements react to form more than one compound, a fixed mass of one element will react with masses of the other element in a ratio of small, whole numbers . For example, copper and chlorine can form a green, crystalline solid with a mass ratio of 0.558 g chlorine to 1 g copper, as well as a brown crystalline solid with a mass ratio of 1.116 g chlorine to 1 g copper. These ratios by themselves may not seem particularly interesting or informative; however, if we take a ratio of these ratios, we obtain a useful and possibly surprising result: a small, whole-number ratio.

\[\mathrm{\dfrac{\dfrac{1.116\: g\: Cl}{1\: g\: Cu}}{\dfrac{0.558\: g\: Cl}{1\: g\: Cu}}=\dfrac{2}{1}}\]

This 2-to-1 ratio means that the brown compound has twice the amount of chlorine per amount of copper as the green compound.

This can be explained by atomic theory if the copper-to-chlorine ratio in the brown compound is 1 copper atom to 2 chlorine atoms, and the ratio in the green compound is 1 copper atom to 1 chlorine atom. The ratio of chlorine atoms (and thus the ratio of their masses) is therefore 2 to 1 (Figure \(\PageIndex{4}\)).

Figure \(\PageIndex{4}\) : Compared to the copper chlorine compound in (a), where copper is represented by brown spheres and chlorine by green spheres, the copper chlorine compound in (b) has twice as many chlorine atoms per copper atom. (credit a: modification of work by “Benjah-bmm27”/Wikimedia Commons; credit b: modification of work by “Walkerma”/Wikimedia Commons)

Example \(\PageIndex{2}\): L aws of Definite and Multiple Proportions

A sample of compound A (a clear, colorless gas) is analyzed and found to contain 4.27 g carbon and 5.69 g oxygen. A sample of compound B (also a clear, colorless gas) is analyzed and found to contain 5.19 g carbon and 13.84 g oxygen. Are these data an example of the law of definite proportions, the law of multiple proportions, or neither? What do these data tell you about substances A and B?

In compound A, the mass ratio of carbon to oxygen is:

\[\mathrm{\dfrac{1.33\: g\: O}{1\: g\: C}} \nonumber\]

In compound B, the mass ratio of carbon to oxygen is:

\[\mathrm{\dfrac{2.67\: g\: O}{1\: g\: C}} \nonumber\]

The ratio of these ratios is:

\[\mathrm{\dfrac{\dfrac{1.33\: g\: O}{1\: g\: C}}{\dfrac{2.67\: g\: O}{1\: g\: C}}=\dfrac{1}{2}} \nonumber\]

This supports the law of multiple proportions. This means that A and B are different compounds, with A having one-half as much carbon per amount of oxygen (or twice as much oxygen per amount of carbon) as B. A possible pair of compounds that would fit this relationship would be A = CO 2 and B = CO.

Exercise \(\PageIndex{2}\)

A sample of compound X (a clear, colorless, combustible liquid with a noticeable odor) is analyzed and found to contain 14.13 g carbon and 2.96 g hydrogen. A sample of compound Y (a clear, colorless, combustible liquid with a noticeable odor that is slightly different from X’s odor) is analyzed and found to contain 19.91 g carbon and 3.34 g hydrogen. Are these data an example of the law of definite proportions, the law of multiple proportions, or neither? What do these data tell you about substances X and Y?

In compound X, the mass ratio of carbon to hydrogen is \(\mathrm{\dfrac{14.13\: g\: C}{2.96\: g\: H}}\).

In compound Y, the mass ratio of carbon to oxygen is \(\mathrm{\dfrac{19.91\: g\: C}{3.34\: g\: H}}\).

The ratio of these ratios is

\[\mathrm{\dfrac{\dfrac{14.13\: g\: C}{2.96\: g\: H}}{\dfrac{19.91\: g\: C}{3.34\: g\: H}}=\dfrac{4.77\: g\: C/g\: H}{5.96\: g\: C/g\: H}=0.800=\dfrac{4}{5}}. \nonumber\]

This small, whole-number ratio supports the law of multiple proportions. This means that X and Y are different compounds.

In the two centuries since Dalton developed his ideas, scientists have made significant progress in furthering our understanding of atomic theory. Much of this came from the results of several seminal experiments that revealed the details of the internal structure of atoms. Here, we will discuss some of those key developments, with an emphasis on application of the scientific method, as well as understanding how the experimental evidence was analyzed. While the historical persons and dates behind these experiments can be quite interesting, it is most important to understand the concepts resulting from their work.

Atomic Theory after the Nineteenth Century

If matter were composed of atoms, what were atoms composed of? Were they the smallest particles, or was there something smaller? In the late 1800s, a number of scientists interested in questions like these investigated the electrical discharges that could be produced in low-pressure gases, with the most significant discovery made by English physicist J. J. Thomson using a cathode ray tube. This apparatus consisted of a sealed glass tube from which almost all the air had been removed; the tube contained two metal electrodes. When high voltage was applied across the electrodes, a visible beam called a cathode ray appeared between them. This beam was deflected toward the positive charge and away from the negative charge, and was produced in the same way with identical properties when different metals were used for the electrodes. In similar experiments, the ray was simultaneously deflected by an applied magnetic field, and measurements of the extent of deflection and the magnetic field strength allowed Thomson to calculate the charge-to-mass ratio of the cathode ray particles. The results of these measurements indicated that these particles were much lighter than atoms (Figure \(\PageIndex{1}\)).

Figure \(\PageIndex{5}\) : (a) J. J. Thomson produced a visible beam in a cathode ray tube. (b) This is an early cathode ray tube, invented in 1897 by Ferdinand Braun. (c) In the cathode ray, the beam (shown in yellow) comes from the cathode and is accelerated past the anode toward a fluorescent scale at the end of the tube. Simultaneous deflections by applied electric and magnetic fields permitted Thomson to calculate the mass-to-charge ratio of the particles composing the cathode ray. (credit a: modification of work by Nobel Foundation; credit b: modification of work by Eugen Nesper; credit c: modification of work by “Kurzon”/Wikimedia Commons).

Based on his observations, here is what Thomson proposed and why: The particles are attracted by positive (+) charges and repelled by negative (−) charges, so they must be negatively charged (like charges repel and unlike charges attract); they are less massive than atoms and indistinguishable, regardless of the source material, so they must be fundamental, subatomic constituents of all atoms. Although controversial at the time, Thomson’s idea was gradually accepted, and his cathode ray particle is what we now call an electron , a negatively charged, subatomic particle with a mass more than one thousand-times less that of an atom. The term “electron” was coined in 1891 by Irish physicist George Stoney, from “ electr ic i on .”

In 1909, more information about the electron was uncovered by American physicist Robert A. Millikan via his “oil drop” experiments. Millikan created microscopic oil droplets, which could be electrically charged by friction as they formed or by using X-rays. These droplets initially fell due to gravity, but their downward progress could be slowed or even reversed by an electric field lower in the apparatus. By adjusting the electric field strength and making careful measurements and appropriate calculations, Millikan was able to determine the charge on individual drops (Figure \(\PageIndex{2}\)).

Figure \(\PageIndex{6}\): Millikan’s experiment measured the charge of individual oil drops. The tabulated data are examples of a few possible values.

Looking at the charge data that Millikan gathered, you may have recognized that the charge of an oil droplet is always a multiple of a specific charge, 1.6 \(\times\) 10 −19 C. Millikan concluded that this value must therefore be a fundamental charge—the charge of a single electron—with his measured charges due to an excess of one electron (1 times 1.6 \(\times\) 10 −19 C), two electrons (2 times 1.6 \(\times\) 10 −19 C), three electrons (3 times 1.6 \(\times\) 10 −19 C), and so on, on a given oil droplet. Since the charge of an electron was now known due to Millikan’s research, and the charge-to-mass ratio was already known due to Thomson’s research (1.759 \(\times\) 10 11 C/kg), it only required a simple calculation to determine the mass of the electron as well.

\[\mathrm{Mass\: of\: electron=1.602\times 10^{-19}\:\cancel{C}\times \dfrac{1\: kg}{1.759\times 10^{11}\:\cancel{C}}=9.107\times 10^{-31}\:kg} \tag{2.3.1}\]

Scientists had now established that the atom was not indivisible as Dalton had believed, and due to the work of Thomson, Millikan, and others, the charge and mass of the negative, subatomic particles—the electrons—were known. However, the positively charged part of an atom was not yet well understood. In 1904, Thomson proposed the “plum pudding” model of atoms, which described a positively charged mass with an equal amount of negative charge in the form of electrons embedded in it, since all atoms are electrically neutral. A competing model had been proposed in 1903 by Hantaro Nagaoka , who postulated a Saturn-like atom, consisting of a positively charged sphere surrounded by a halo of electrons (Figure \(\PageIndex{3}\)).

Figure \(\PageIndex{7}\) : (a) Thomson suggested that atoms resembled plum pudding, an English dessert consisting of moist cake with embedded raisins (“plums”). (b) Nagaoka proposed that atoms resembled the planet Saturn, with a ring of electrons surrounding a positive “planet.” (credit a: modification of work by “Man vyi”/Wikimedia Commons; credit b: modification of work by “NASA”/Wikimedia Commons).

The next major development in understanding the atom came from Ernest Rutherford , a physicist from New Zealand who largely spent his scientific career in Canada and England. He performed a series of experiments using a beam of high-speed, positively charged alpha particles (α particles) that were produced by the radioactive decay of radium; α particles consist of two protons and two neutrons (you will learn more about radioactive decay in the chapter on nuclear chemistry). Rutherford and his colleagues Hans Geiger (later famous for the Geiger counter) and Ernest Marsden aimed a beam of α particles, the source of which was embedded in a lead block to absorb most of the radiation, at a very thin piece of gold foil and examined the resultant scattering of the α particles using a luminescent screen that glowed briefly where hit by an α particle.

What did they discover? Most particles passed right through the foil without being deflected at all. However, some were diverted slightly, and a very small number were deflected almost straight back toward the source (Figure \(\PageIndex{4}\)). Rutherford described finding these results: “It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you” 1 (p. 68).

Figure \(\PageIndex{8}\): Geiger and Rutherford fired α particles at a piece of gold foil and detected where those particles went, as shown in this schematic diagram of their experiment. Most of the particles passed straight through the foil, but a few were deflected slightly and a very small number were significantly deflected.

Here is what Rutherford deduced: Because most of the fast-moving α particles passed through the gold atoms undeflected, they must have traveled through essentially empty space inside the atom. Alpha particles are positively charged, so deflections arose when they encountered another positive charge (like charges repel each other). Since like charges repel one another, the few positively charged α particles that changed paths abruptly must have hit, or closely approached, another body that also had a highly concentrated, positive charge. Since the deflections occurred a small fraction of the time, this charge only occupied a small amount of the space in the gold foil. Analyzing a series of such experiments in detail, Rutherford drew two conclusions:

• The volume occupied by an atom must consist of a large amount of empty space.
• A small, relatively heavy, positively charged body, the nucleus , must be at the center of each atom.

This analysis led Rutherford to propose a model in which an atom consists of a very small, positively charged nucleus, in which most of the mass of the atom is concentrated, surrounded by the negatively charged electrons, so that the atom is electrically neutral (Figure \(\PageIndex{5}\)).

Figure \(\PageIndex{9}\) : The α particles are deflected only when they collide with or pass close to the much heavier, positively charged gold nucleus. Because the nucleus is very small compared to the size of an atom, very few α particles are deflected. Most pass through the relatively large region occupied by electrons, which are too light to deflect the rapidly moving particles.

After many more experiments, Rutherford also discovered that the nuclei of other elements contain the hydrogen nucleus as a “building block,” and he named this more fundamental particle the proton , the positively charged, subatomic particle found in the nucleus. With one addition, which you will learn next, this nuclear model of the atom, proposed over a century ago, is still used today.

Another important finding was the discovery of isotopes. During the early 1900s, scientists identified several substances that appeared to be new elements, isolating them from radioactive ores. For example, a “new element” produced by the radioactive decay of thorium was initially given the name mesothorium. However, a more detailed analysis showed that mesothorium was chemically identical to radium (another decay product), despite having a different atomic mass. This result, along with similar findings for other elements, led the English chemist Frederick Soddy to realize that an element could have types of atoms with different masses that were chemically indistinguishable. These different types are called isotopes —atoms of the same element that differ in mass. Soddy was awarded the Nobel Prize in Chemistry in 1921 for this discovery.

One puzzle remained: The nucleus was known to contain almost all of the mass of an atom, with the number of protons only providing half, or less, of that mass. Different proposals were made to explain what constituted the remaining mass, including the existence of neutral particles in the nucleus. As you might expect, detecting uncharged particles is very challenging, and it was not until 1932 that James Chadwick found evidence of neutrons , uncharged, subatomic particles with a mass approximately the same as that of protons. The existence of the neutron also explained isotopes: They differ in mass because they have different numbers of neutrons, but they are chemically identical because they have the same number of protons. This will be explained in more detail later in this unit.

Video \(\PageIndex{2}\): An Introduction to Subatomic Particles

Video \(\PageIndex{3}\): A summary of discoveries in atomic theory.

Video \(\PageIndex{4}\): A different summary of discoveries in atomic theory.

The ancient Greeks proposed that matter consists of extremely small particles called atoms. Dalton postulated that each element has a characteristic type of atom that differs in properties from atoms of all other elements, and that atoms of different elements can combine in fixed, small, whole-number ratios to form compounds. Samples of a particular compound all have the same elemental proportions by mass. When two elements form different compounds, a given mass of one element will combine with masses of the other element in a small, whole-number ratio. During any chemical change, atoms are neither created nor destroyed.

Although no one has actually seen the inside of an atom, experiments have demonstrated much about atomic structure. Thomson’s cathode ray tube showed that atoms contain small, negatively charged particles called electrons. Millikan discovered that there is a fundamental electric charge—the charge of an electron. Rutherford’s gold foil experiment showed that atoms have a small, dense, positively charged nucleus; the positively charged particles within the nucleus are called protons. Chadwick discovered that the nucleus also contains neutral particles called neutrons. Soddy demonstrated that atoms of the same element can differ in mass; these are called isotopes.

• Ernest Rutherford, “The Development of the Theory of Atomic Structure,” ed. J. A. Ratcliffe, in Background to Modern Science , eds. Joseph Needham and Walter Pagel, (Cambridge, UK: Cambridge University Press, 1938), 61–74. Accessed September 22, 2014, https://ia600508.us.archive.org/3/it...e032734mbp.pdf .

Contributors

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] ).

• Adelaide Clark, Oregon Institute of Technology
• Crash Course Chemistry: Crash Course is a division of Complexly and videos are free to stream for educational purposes.
• TED-Ed’s commitment to creating lessons worth sharing is an extension of TED’s mission of spreading great ideas. Within TED-Ed’s growing library of TED-Ed animations, you will find carefully curated educational videos, many of which represent collaborations between talented educators and animators nominated through the TED-Ed website .

Have feedback to give about this text? Click here .

Found a typo and want extra credit? Click here .