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What is the model of the atom proposed by Ernest Rutherford?

What is the rutherford gold-foil experiment, what were the results of rutherford's experiment, what did ernest rutherford's atomic model get right and wrong, what was the impact of ernest rutherford's theory.

Rutherford model

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• UC Davis - The Rutherford Scattering Experiment
• Chemistry LibreTexts - Rutherford's Experiment- The Nuclear Model of the Atom

The atom , as described by Ernest Rutherford , has a tiny, massive core called the nucleus . The nucleus has a positive charge. Electrons are particles with a negative charge. Electrons orbit the nucleus. The empty space between the nucleus and the electrons takes up most of the volume of the atom.

A piece of gold foil was hit with alpha particles , which have a positive charge. Most alpha particles went right through. This showed that the gold atoms were mostly empty space. Some particles had their paths bent at large angles. A few even bounced backward. The only way this would happen was if the atom had a small, heavy region of positive charge inside it.

The previous model of the atom, the Thomson atomic model , or the “plum pudding” model, in which negatively charged electrons were like the plums in the atom’s positively charged pudding, was disproved. The Rutherford atomic model relied on classical physics. The Bohr atomic model , relying on quantum mechanics, built upon the Rutherford model to explain the orbits of electrons.

The Rutherford atomic model was correct in that the atom is mostly empty space. Most of the mass is in the nucleus, and the nucleus is positively charged. Far from the nucleus are the negatively charged electrons. But the Rutherford atomic model used classical physics and not quantum mechanics. This meant that an electron circling the nucleus would give off electromagnetic radiation . The electron would lose energy and fall into the nucleus. In the Bohr model, which used quantum theory, the electrons exist only in specific orbits and can move between these orbits.​

The gold-foil experiment showed that the atom consists of a small, massive, positively charged nucleus with the negatively charged electrons being at a great distance from the centre. Niels Bohr built upon Rutherford’s model to make his own. In Bohr’s model the orbits of the electrons were explained by quantum mechanics.

Rutherford model , description of the structure of atoms proposed (1911) by the New Zealand-born physicist Ernest Rutherford . The model described the atom as a tiny, dense, positively charged core called a nucleus, in which nearly all the mass is concentrated, around which the light, negative constituents , called electrons , circulate at some distance, much like planets revolving around the Sun .

The nucleus was postulated as small and dense to account for the scattering of alpha particles from thin gold foil, as observed in a series of experiments performed by undergraduate Ernest Marsden under the direction of Rutherford and German physicist Hans Geiger in 1909. A radioactive source emitting alpha particles (i.e., positively charged particles, identical to the helium atom nucleus and 7,000 times more massive than electrons) was enclosed within a protective lead shield. The radiation was focused into a narrow beam after passing through a slit in a lead screen. A thin section of gold foil was placed in front of the slit, and a screen coated with zinc sulfide to render it fluorescent served as a counter to detect alpha particles. As each alpha particle struck the fluorescent screen , it produced a burst of light called a scintillation, which was visible through a viewing microscope attached to the back of the screen. The screen itself was movable, allowing Rutherford and his associates to determine whether or not any alpha particles were being deflected by the gold foil.

Most alpha particles passed straight through the gold foil, which implied that atoms are mostly composed of open space. Some alpha particles were deflected slightly, suggesting interactions with other positively charged particles within the atom. Still other alpha particles were scattered at large angles, while a very few even bounced back toward the source. (Rutherford famously said later, “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.”) Only a positively charged and relatively heavy target particle, such as the proposed nucleus, could account for such strong repulsion. The negative electrons that balanced electrically the positive nuclear charge were regarded as traveling in circular orbits about the nucleus. The electrostatic force of attraction between electrons and nucleus was likened to the gravitational force of attraction between the revolving planets and the Sun. Most of this planetary atom was open space and offered no resistance to the passage of the alpha particles.

The Rutherford model supplanted the “plum-pudding” atomic model of English physicist Sir J.J. Thomson , in which the electrons were embedded in a positively charged atom like plums in a pudding. Based wholly on classical physics , the Rutherford model itself was superseded in a few years by the Bohr atomic model , which incorporated some early quantum theory . See also atomic model .

May, 1911: Rutherford and the Discovery of the Atomic Nucleus

In 1909, Ernest Rutherford’s student reported some unexpected results from an experiment Rutherford had assigned him. Rutherford called this news the most incredible event of his life.

In the now well-known experiment, alpha particles were observed to scatter backwards from a gold foil. Rutherford’s explanation, which he published in May 1911, was that the scattering was caused by a hard, dense core at the center of the atom–the nucleus.

Ernest Rutherford was born in New Zealand, in 1871, one of 12 children. Growing up, he often helped out on the family farm, but he was a good student, and received a scholarship to attend the University of New Zealand. After college he won a scholarship in 1894 to become a research student at Cambridge. Upon receiving the news of this scholarship, Rutherford is reported to have said, “That’s the last potato I’ll ever dig.”

At Cambridge, the young Rutherford worked in the Cavendish lab with J.J. Thomson, discoverer of the electron. Rutherford’s talent was quickly recognized, and in 1898 he took a professorship at McGill University in Montreal. There, he identified alpha and beta radiation as two separate types of radiation, and studied some of their properties, though he didn’t know that alphas were helium nuclei. In 1901 Rutherford and chemist Frederick Soddy found that one radioactive element can decay into another. The discovery earned Rutherford the 1908 Nobel Prize in Chemistry, which irritated him somewhat because he considered himself a physicist, not a chemist. (Rutherford is widely quoted as having said, “All science is either physics or stamp collecting”)

In 1907 Rutherford returned to England, to the University of Manchester. In 1909, he and his colleague Hans Geiger were looking for a research project for a student, Ernest Marsden. Rutherford had already been studying the scattering of alpha particles off a gold target, carefully measuring the small forward angles through which most of the particles scattered. Rutherford, who didn’t want to neglect any angle of an experiment, no matter how unpromising, suggested Marsden look to see if any alpha particles actually scattered backwards.

Marsden was not expected to find anything, but nonetheless he dutifully and carefully carried out the experiment. He later wrote that he felt it was a sort of test of his experimental skills. The experiment involved firing alpha particles from a radioactive source at a thin gold foil. Any scattered particles would hit a screen coated with zinc sulfide, which scintillates when hit with charged particles. Marsden was to sit in the darkened room, wait for his eyes to adjust to the darkness, and then patiently stare at the screen, expecting to see nothing at all.

Instead, Marsden saw lots of tiny, fleeting flashes of yellowish light, on average more than one blip per second.

He could hardly believe what he saw. He tested and retested every aspect of the experiment, but when he couldn’t find anything wrong, he reported the results to Rutherford.

Rutherford too was astonished. As he was fond of saying, “It was as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you."

About one in every few thousand of the alpha particles fired at the gold target had scattered at an angle greater than 90 degrees. This didn’t fit with the prevailing model of the atom, the so-called plum pudding model developed by J.J. Thomson. In this model electrons were believed to be stuck throughout a blob of positively charged matter, like raisins in a pudding. But this sort of arrangement would only cause small angle scattering, nothing like what Marsden had observed.

After thinking about the problem for over a year, Rutherford came up with an answer. The only explanation, Rutherford suggested in 1911, was that the alpha particles were being scattered by a large amount of positive charge concentrated in a very small space at the center of the gold atom. The electrons in the atom must be orbiting around this central core, like planets around the sun, Rutherford proposed.

Rutherford carried out a fairly simple calculation to find the size of the nucleus, and found it to be only about 1/100,000 the size of the atom. The atom was mostly empty space.

In March 1911, Rutherford announced his surprising finding at a meeting of the Manchester Literary and Philosophical Society, and in May 1911, he published a paper on the results in the Philosophical Magazine .

Later Rutherford and Marsden tried the experiment with other elements as the target, and measured their nuclei as well.

The solar system model was not immediately accepted. One obvious problem was that according to Maxwell’s equations, electrons traveling in a circular orbit should radiate energy, and therefore slow down and fall into the nucleus. A solar system atom wouldn’t last long.

Fortunately, Niels Bohr was soon able to save the solar system model by applying new ideas from quantum mechanics. He showed that the atom could stay intact if electrons were only allowed to occupy certain discrete orbitals.

Though Rutherford still didn’t know what was in this nucleus he had discovered (protons and neutrons would be identified later), his insight in 1911, which overturned the prevailing plum pudding model of the atom, had opened the way for modern nuclear physics.

Ernie Tretkoff

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Alpha Particles and the Atom

Rutherford at manchester, 1907–1919.

Ernest Rutherford discovered the nucleus of the atom in 1911. We read this in textbooks and in popular writings. But what does that statement mean? Geographical discovery usually means that one sees a place for the first time. But can discovery be the same for a realm hidden from sight? One cannot see an atom in that sense. So this hints that perhaps the story of the discovery of the nucleus was more complicated. The story as it unfolded in Rutherford's lab at the University in Manchester revolved around real people. It involved frustrations and triumphs. It involved hard work and perplexity and inspiration.

Rutherford arrived in Manchester in the summer of 1907, months before the university's term began. He had been named Langworthy Professor of Physics, successor to Arthur Schuster (1851–1934), who retired at age 56 to recruit Rutherford. Schuster had built a modern physics building, hired Hans Geiger, Ph.D. (1882–1945) because of his experimental skill, and endowed a new position in mathematical physics to round out a full physics program. Rutherford entered the center of the physics world. Researchers came to him by the dozen.

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I found Rutherford's place very busy, hard working. But a very dirty place. Namely, Manchester is very foggy, foggy and smoky. And of course everywhere you see smoke there, everywhere the smoke. Now the technique used in Rutherford’s lab was to fit up an electroscope. You have to build it yourself of cocoa boxes, gold leaf and sulfur isolation. And you charge the electroscope by sealing wax which you rubbed on your trousers. So it was a very primitive technique. But of course also a microscope to read the electroscope. Now the microscope was fixed and then you were not supposed to touch it. And of course you were not supposed to clean it. So years went on without apparatus being cleaned. But apart from the shortcomings it was a very fine lab, nice rooms, and lots of people working there—able people.... I remember Moseley very well, with whom I was on very friendly terms. I will tell you later about his work. And Charles Darwin was there. He was lecturing in theoretical physics. And Russell, who later came to Oxford. An Italian, Rossi, did spectroscopic work. He showed that ionium and sodium have the same spectrum. And then Geiger was there. He was an assistant. And also an assistant named Makower, who died since. Geiger and Makower published a book together. And also a chap Robinson, who worked on beta rays. Gray, a New Zealand man. Marsden who came from Australia. Fajans who came from Germany. And Boltwood was there for a while. He came from Yale. Rutherford invited him in hope that Boltwood, a great chemist, would purify ionium, but he failed as many others.

Rutherford arrived with many research questions in mind. He was not done with the puzzles of the decay families of thorium, radium, etc., but he was passing much of this work to Boltwood, Hahn, and Soddy. Boltwood and Hahn both worked with Rutherford in Manchester, Boltwood in 1909–1910 and Hahn in 1907–1908. Rutherford was gradually turning his attention much more to the α (alpha), β (beta), and γ (gamma) rays themselves and to what they might reveal about the atom. That is, he was leaving radio-chemistry to others and turning to physics.

Rutherford's early team at Manchester included Geiger and William Kay (1879–1961), junior laboratory assistant since 1894. Rutherford promoted Kay to laboratory steward in 1908, to manage lab equipment and to aid him in his research. In 1957, Kay thought back to his youth with Rutherford in an interview. The language is quaint, but the description is as close to Rutherford's approach as we get. The questioner was Samuel Devons (1914–2006), who was one of Rutherford's last students in the 1930s.

[Devons] “When you were here [in Manchester], during this period... did Rutherford actually make any apparatus himself?” [Kay] “No, no, no, no. We used to, I used to set up nearly all his apparatus. You know, when he did his work, you know, oftener than not, he used to tell me and we did a rough experiment, re...” [D.] “Did he sketch out what he wanted?” [K.] “Well, he'd tell you what he wanted, roughly, you see, but he'd let you make what you wanted, you see, he'd tell you what he was going to do, which was very good, you see. It gives you......... it learnt you a lot and you knew what to do and what not to do. And then we would do a rough experiment, and get one or two curves you see, and then straight away button it on to somebody else to do the real work, and that's how he did his........ attacked these little things, you see.” [D.] “He tried them out himself first?” [K.] “He'd try a rough experiment himself on the little things, d'you see, and then he'd turn it over on to somebody...” (Quoted in Hughes, p. 104)

Rutherford and Hans Geiger worked closely in 1907 and 1908 on the detection and measurement of α particles. If they were to use α particles to probe the atom, they had first to know more about these particles and their behavior. Rutherford had tried and failed back at McGill to count α particles.

A year later in Manchester, he and Geiger succeeded with two methods of observing α particles. The first method involved scintillations excited by α particles on a thin layer of zinc sulfide. They observed these through a microscope and counted the scintillations at different angles of dispersion. They also developed an "electrometer" that could demonstrate the passage of an individual α particle to a large audience. The instrument, which evolved into the "Geiger counter," had a partially evacuated metal cylinder with a wire down its center. They applied a voltage between the cylinder and the wire high enough almost to spark. They admitted α particles through a thin mica window, where these particles collided with gasses, producing gas ions. These then collided with other molecules and produced more ions, and so on. Each α particle produced a cascade of ions, which partially discharged the cylinder and indicated the passage of an α particle. Geiger and Rutherford published several articles in 1908 and 1909 on these methods and their use.

Rutherford wrote to Henry Bumstead (1870–1920), an American physicist, on 11 July 1908:

Geiger is a good man and worked like a slave. I could never have found time for the drudgery before we got things going in good style. Finally all went well, but the scattering is the devil. Our tube worked like a charm and we could easily get a throw of 50 mm. for each particle. ... Geiger is a demon at the work of counting scintillations and could count at intervals for a whole night without disturbing his equanimity. I damned vigorously and retired after two minutes. (Quoted in Eve, p. 180.)

Although Rutherford suspected as early as 1906 that α particles were helium atoms stripped of their electrons, he demanded a high standard of proof. One kind of experiment was not enough. One kind of detector was not enough. He wanted more proof. For this, Rutherford desired "big voltages" and big electromagnets to divert α particles, but this method was not yet ripe. Lab steward William Kay recalled in the cited oral history interview that Rutherford in 1908 insisted that strong electric and magnetic fields were needed to measure more directly the charge and mass of the α and β particles:

And that's what he was after all the time. That's what he got at Cambridge [after 1919], which we never got here, you see, because we'd got no money. (Hughes, “William Kay,” 2008, pp. 109–110.)

Kay said Rutherford wanted a big, water-cooled magnet, but that he “dropped it like a hot cake” when he learned its cost. So he needed a new line of attack. The new line was very simple, a chemical procedure mixed with physics. For this work Rutherford recruited Thomas Royds (1884–1955), who had earned his Physics Honours degree in 1906. They collected α particles in a sealed glass tube, compressed them, and passed an electric spark through. They studied the emitted light in a spectroscope and found it to be identical to the spectrum of helium. Within a few months, Rutherford was awarded the Nobel Prize for Chemistry, "for his investigations into the disintegration of the elements, and the chemistry of radioactive substances." (Nobel citation) Rutherford and Royds had established the identity and primary properties of α particles. Rutherford next turned his attention to using them to probe the atom.

The autumn of 1908 began an important series of researches. Geiger had been passing beams of α particles through gold and other metallic foils, using the new detection techniques to measure how much these beams were dispersed by the atoms in the foils. Geiger thought Ernest Marsden (1889–1970), a 19-year-old student in Honours Physics, was ready to help on these experiments and suggested it to Rutherford. Since Rutherford often pushed third-year students into research, saying this was the best way to learn about physics, he readily agreed.

Geiger and Marsden began with small-angle dispersion and tried various thicknesses of foils, seeking mathematical relationships between dispersion and thickness of foil or number of atoms traversed. Marsden later recalled that Rutherford said to him amidst these experiments: "See if you can get some effect of alpha-particles directly reflected from a metal surface." (Reported by Marsden in Birks, 1962, p. 8). Marsden doubted that Rutherford expected back scatter of α particles, but as Marsden wrote

...it was one of those 'hunches' that perhaps some effect might be observed, and that in any case that neighbouring territory of this Tom Tiddler's ground might be explored by reconnaissance. Rutherford was ever ready to meet the unexpected and exploit it, where favourable, but he also knew when to stop on such excursions. (Birks, 1962, p. 8)

This was Rutherford's playful approach in action. His students and others tried out his ideas, many of which were dead-ends. This idea to look for backscattering of α particles, however, paid off. Rutherford wrote:

Experiment, directed by the disciplined imagination either of an individual or, still better, of a group of individuals of varied mental outlook, is able to achieve results which far transcend the imagination alone of the greatest philosopher. (Quoted in Eve, 1939, Frontmatter)

Sometime later in 1908 or 1909, Marsden said, he reported his results to Rutherford. Rutherford recalled this a little differently:

I remember ...later Geiger coming to me in great excitement and saying, 'We have been able to get some of the α -particles coming backwards...' It was quite the most incredible event that has ever happened to me in my life. It was almost incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you. ( Rutherford , 1938, p. 68)

Human memory is fallible. Whether Marsden or Geiger told Rutherford, the effect was the same. Rutherford said they should prepare a publication from this research, which they submitted in May 1909. Moreover, this started Rutherford thinking toward what ultimately, almost two years later, he published as a theory of the atom.

What was Rutherford doing for the rest of 1909 and all of 1910? For one thing, his close friend Boltwood was in Manchester for the academic year working with Rutherford on radioactive decay products of radium. He was also reviewing and speaking on earlier ideas about atomic structure. Most importantly, he was taking the phenomenon of the scattering of α particles apart systematically and testing each piece. Rutherford did not have his bold idea — the nuclear atom — instantly, but he came to it gradually by considering the problem from many sides.

In the autumn of 1910 he brought Marsden back to Manchester to complete rigorous experimental testing of his ideas with Geiger. They re-established rates of emission and the ranges of α particles by radioactive sources and they re-examined their statistical analyses. Rutherford tried to reconcile scattering results with different atomic models, especially that of J.J. Thomson, in which the positive electricity was considered as dispersed evenly throughout the whole sphere of the atom.

At some point in the winter of 1910–1911, Rutherford worked out the basic idea of an atom with a "charged center." As Geiger and Marsden pointed out in their 1909 article:

If the high velocity and mass of the α -particle be taken into account, it seems surprising that some of the α -particles, as the experiment shows, can be turned within a layer of 6 x 10 -5 cm. of gold through an angle of 90°, and even more. To produce a similar effect by a magnetic field, the enormous field of 109 absolute units would be required. (Birks, p. 179)

Rutherford concluded in his May 1911 paper that such a remarkable deviation in the path of a massive charged particle could only be achieved if most of the mass of, say, an atom of gold and most of its charge were concentrated in a very small central body. Note: at this point in 1911, Rutherford did not call this a "nucleus."

The first public announcement of the nuclear theory by Rutherford was made at a meeting of the Manchester Literary and Philosophical Society, and he invited us young boys to go to the meeting. He said he’d got some interesting things to say and he thought we’d like to hear them. We didn’t know what it was about at that time. The older people in the laboratory did, of course Geiger and Marsden knew because they were already doing the experiments. In fact, unless they had done some which were sufficient to be decisive, Rutherford never mentioned it publicly. And, of course, Darwin knew about it much earlier. But that must have been early in 1911, and we went to the meeting and he told us. And he mentioned then that there was some experimental evidence which had been obtained by Geiger and Marsden. He did not, as far as I remember, say more about the results than that they were quite decisive. And, as I said before, he would never have made a public announcement of that kind if he hadn’t had good evidence. And that is one of the characteristics that runs through all Rutherford’s work, particularly all his work up to the end of the Manchester period. If you look at some of his papers in the early days — I call McGill the early days — he was quite convinced that the alpha particles were atoms of helium, but he never said that in those words. He always said they were either atoms of helium or molecules of hydrogen or perhaps he may have said something else of that weight. It was quite characteristic of him that he would never say a thing was so unless he had experimental evidence for it that really satisfied him.

In fact, Rutherford was exceedingly cautious in drawing conclusions about this central charge: “A simple calculation shows that the atom must be a seat of an intense electric field in order to produce such a large deflexion at a single encounter.” (Birks, p. 183). He worked out quickly and roughly that several quantitative relationships should be true if this basic theory were correct. First, the number of α particles scattered through a given angle should be proportional to the thickness of the foil. Second, that number should be proportional to the square of the nuclear charge. Lastly, it should be inversely proportional to the fourth power of the velocity of the α particle. These three ideas laid out the experimental program of Geiger and Marsden for the next year.

Rutherford’s interest was then almost entirely in the research. He had done very little teaching in McGill. He was research professor. I suppose he gave some lectures but it would have been very few. And his interest was quite naturally on the research side. He did give some lectures, but elementary lectures, the kind of thing you would expect a man to know before he came to the University. They were the lectures to the engineers. They were a rowdy lot and Rutherford could keep them under control. There was perhaps only one other man in the department who could have done it, and he (Rutherford?) enjoyed them because he was able to show them the very interesting experiments one can perform in elementary courses.

It's often been said to me that Rutherford was a bad lecturer. I never heard such nonsense. It is quite true that on occasion he would be a bit dull, a bit mixed up, but that was only on very rare occasions. There were other occasions when he was really most stimulating. There was a tremendous enthusiasm about him.

Rutherford entertained the possibility that the charged center is negative. That sounds odd today, so what made it reasonable? First, it wasn't very different from Thomson's model. Second, since Rutherford knew that α particles carry a double + charge, he thought this might act the same way the Sun does on a comet sweeping near it. It would slingshot the α particle around and back towards its source. He also considered a nearly forgotten model suggested by Japanese physicist Hantaro Nagaoka (1865–1950) — the Saturnian model. Nagaoka and Rutherford were in contact in 1910 and 1911 and Rutherford mentioned Nagaoka's model of "a central attracting mass surround by rings of rotating electrons" (Birks, p. 203). The end result in this critical Rutherford paper, however, was Rutherford's announcement that whether the atom were a disk or a sphere, and indeed whether the central charge were positive or negative, would not affect the calculations. Rutherford was always careful not to claim more than his results could support.

Rutherford did see possible tests of the nature of the central charge. The absorption of β particles, he said, should be different with a negative center versus a positive one. A positive center would explain the great velocity that α particles achieve during emission from radioactive elements. But these were only hints.

Geiger and Marsden did indeed work systematically through the testable implications of Rutherford's central charge hypothesis. The first major publication of their results was in German in the Proceedings of the Vienna Academy of Sciences ( Sitzungberichte der Wiener Akademie der Wissenschaften) in 1912. This 30-page version was followed by one in English in 1913 in the Philosophical Magazine: "The Laws of Deflexion of α Particles through Large Angles" The English version is the better known. Slight differences between the two led one historian to suggest that Rutherford decided in favor of a positively charged center by August 1912 (Trenn, 1974). Rutherford's other team members, especially Charles Galton Darwin (1887–1962), H.G.J. Moseley (1887–1915), and Niels Bohr (1885–1962) figured prominently in the ultimate establishment of Rutherford's nuclear atom.

The ‘Great War’ totally disrupted work in Rutherford's Manchester department. Bohr returned to Denmark. Marsden accepted a professorship in New Zealand. Moseley died in the Battle of Gallipoli. James Chadwick (1891–1974), who was working with Geiger at the Technical University of Berlin when war broke out, spent several years interned in the Ruhleben camp for prisoners of war. Other students went off to war, too, and Rutherford devoted considerable energy to mobilizing science for the war effort and specifically to anti-submarine techniques.

Against this distracted background, Rutherford and his lab steward, William Kay, began in 1917 to explore the passage of α particles through hydrogen, nitrogen, and other gases. When the Great War ended, Ernest Marsden briefly helped with the tedious scintillation observations that provided clues to the nature of the nucleus. Rutherford reported the tentative results of these extensive experiments in 1919. Rutherford placed a source of radium C (bismuth-214) in a sealable brass container, fitted so that the position of the source could be changed and so that different gases could be introduced or a vacuum produced, as desired. The α particles traversed the interior of the container and passed through a slit, covered by a silver plate or other material, and hit a zinc sulfide screen, where a scintillation was observed in a darkened room. When hydrogen gas was introduced into the container and care was taken to absorb the α particles before they hit the screen, scintillations were still observed. Rutherford posited that as the α particles traversed the hydrogen gas, they occasionally collided with hydrogen nuclei. As Rutherford wrote, this produced “swift hydrogen atoms” which were mostly projected forward in the direction of the α particles’ original motion.

Rutherford had several subtle questions in mind during these experiments, mostly concerned with the nature of the nucleus. He asked his colleague Darwin to analyze these collisions based on a ‘simple theory’ of elastic collisions between point nuclei repelled according to an inverse square law, the α particles carrying a charge of 2 times that of an electron (and of opposite sign) and the hydrogen nuclei 1 times. Darwin found that all α particles approaching within 2.4x10 -13 cm would produce a ‘swift hydrogen atom.’ This simple theory, however, predicted far fewer accelerated hydrogen atoms than were observed in the experiments.

Rutherford rejected explanations of this variance based on different charges on the particles or other laws than inverse square laws. Rather, he concluded that for distances on the order of the diameter of the electron, ‘the structure of the helium nucleus can no longer be regarded as a point…’. He posited that the helium nucleus ( α particle) has a complex structure of four hydrogen nuclei plus two negatively charged electrons. (We would say it is composed of two protons.) Rutherford concluded that deformation of complex nuclei during collisions was a more likely explanation, the variation of the forces between the nuclei varying in a complex way on close approach.

Taking into account the intense forces brought into play in such collisions, it would not be surprising if the helium nucleus were to break up. No evidence of such a disintegration…has been observed, indicating that the helium nucleus must be a very stable structure.

We must remember that Rutherford could not directly observe the structure of the nucleus, so his conclusions were tentative. Nevertheless, he was openly considering the possibilities of a complex nucleus, capable of deformation and even of possible disintegration. These thoughts shaped this intense period of experimental researches.

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Experimental Evidence for the Structure of the Atom

George sivulka march 23, 2017, submitted as coursework for ph241 , stanford university, winter 2017, introduction.

The Rutherford Gold Foil Experiment offered the first experimental evidence that led to the discovery of the nucleus of the atom as a small, dense, and positively charged atomic core. Also known as the Geiger-Marsden Experiments, the discovery actually involved a series of experiments performed by Hans Geiger and Ernest Marsden under Ernest Rutherford. With Geiger and Marsden's experimental evidence, Rutherford deduced a model of the atom, discovering the atomic nucleus. His "Rutherford Model", outlining a tiny positively charged atomic center surrounded by orbiting electrons, was a pivotal scientific discovery revealing the structure of the atoms that comprise all the matter in the universe.

The experimental evidence behind the discovery involved the scattering of a particle beam after passing through a thin gold foil obstruction. The particles used for the experiment - alpha particles - are positive, dense, and can be emitted by a radioactive source. Ernest Rutherford discovered the alpha particle as a positive radioactive emission in 1899, and deduced its charge and mass properties in 1913 by analyzing the charge it induced in the air around it. [1] As these alpha particles have a significant positive charge, any significant potential interference would have to be caused by a large concentration of electrostatic force somewhere in the structure of the atom. [2]

Previous Model of the Atom

The scattering of an alpha particle beam should have been impossible according to the accepted model of the atom at the time. This model, outlined by Lord Kelvin and expanded upon by J. J. Thompson following his discovery of the electron, held that atoms were comprised of a sphere of positive electric charge dotted by the presence of negatively charged electrons. [3] Describing an atomic model similar to "plum pudding," it was assumed that electrons were distributed throughout this positive charge field, like plums distributed in the dessert. However, this plum pudding model lacked the presence of any significant concentration of electromagnetic force that could tangibly affect any alpha particles passing through atoms. As such, alpha particles should show no signs of scattering when passing through thin matter. [4] (see Fig. 2)

The Geiger Marsden Experiments

Testing this accepted theory, Hans Geiger and Ernest Marsden discovered that atoms indeed scattered alpha particles, a experimental result completely contrary to Thompson's model of the atom. In 1908, the first paper of the series of experiments was published, outlining the apparatus used to determine this scattering and the scattering results at small angles. Geiger constructed a two meter long glass tube, capped off on one end by radium source of alpha particles and on the other end by a phosphorescent screen that emitted light when hit by a particle. (see Fig. 3) Alpha particles traveled down the length of the tube, through a slit in the middle and hit the screen detector, producing scintillations of light that marked their point of incidence. Geiger noted that "in a good vacuum, hardly and scintillations were observed outside of the geometric image of the slit, "while when the slit was covered by gold leaf, the area of the observed scintillations was much broader and "the difference in distribution could be noted with the naked eye." [5]

On Rutherford's request, Geiger and Marsden continued to test for scattering at larger angles and under different experimental parameters, collecting the data that enabled Rutherford to further his own conclusions about the nature of the nucleus. By 1909, Geiger and Marsden showed the reflection of alpha particles at angles greater than 90 degrees by angling the alpha particle source towards a foil sheet reflector that then would theoretically reflect incident particles at the detection screen. Separating the particle source and the detector screen by a lead barrier to reduce stray emission, they noted that 1 in every 8000 alpha particles indeed reflected at the obtuse angles required by the reflection of metal sheet and onto the screen on the other side. [6] Moreover, in 1910, Geiger improved the design of his first vacuum tube experiment, making it easier to measure deflection distance, vary foil types and thicknesses, and adjust the alpha particle stream' velocity with mica and aluminum obstructions. Here he discovered that both thicker foil and foils made of elements of increased atomic weight resulted in an increased most probable scattering angle. Additionally, he confirmed that the probability for an angle of reflection greater than 90 degrees was "vanishingly small" and noted that increased particle velocity decreased the most probably scattering angle. [7]

Rutherford's Atom

Backed by this experimental evidence, Rutherford outlined his model of the atom's structure, reasoning that as atoms clearly scattered incident alpha particles, the structure contained a much larger electrostatic force than earlier anticipated; as large angle scattering was a rare occurrence, the electrostatic charge source was only contained within a fraction of the total volume of the atom. As he concludes this reasoning with the "simplest explanation" in his 1911 paper, the "atom contains a central charge distributed through a very small volume" and "the large single deflexions are due to the central charge as a whole." In fact, he mathematically modeled the scattering patterns predicted by this model with this small central "nucleus" to be a point charge. Geiger and Marsden later experimentally verified each of the relationships predicted in Rutherford's mathematical model with techniques and scattering apparatuses that improved upon their prior work, confirming Rutherford's atomic structure. [4, 8, 9] (see Fig. 1)

With the experimentally analyzed nature of deflection of alpha rays by thin gold foil, the truth outlining the structure of the atom falls into place. Though later slightly corrected by Quantum Mechanics effects, the understanding of the structure of the the atom today almost entirely follows form Rutherford's conclusions on the Geiger and Marsden experiments. This landmark discovery fundamentally furthered all fields of science, forever changing mankind's understanding of the world around us.

© George Sivulka. The author grants permission to copy, distribute and display this work in unaltered form, with attribution to the author, for noncommercial purposes only. All other rights, including commercial rights, are reserved to the author.

[1] E. Rutherford, "Uranium Radiation and the Electrical Conduction Produced By It," Philos. Mag. 47 , 109 (1899).

[2] E. Rutherford, "The Structure of the Atom," Philos. Mag. 27 , 488 (1914).

[3] J. J. Thomson, "On the Structure of the Atom: an Investigation of the Stability and Periods of Oscillation of a Number of Corpuscles Arranged at Equal Intervals Around the Circumference of a Circle; with Application of the Results to the Theory of Atomic Structure," Philos. Mag. 7 , 237 (1904).

[4] E. Rutherford, "The Scattering of α and β Particles by Matter and the Structure of the Atom," Philos. Mag. 21 , 669 (1911).

[5] H. Geiger, "On the Scattering of the α Particles by Matter," Proc. R. Soc. A 81 , 174 (1908).

[6] H. Geiger and E. Marsden, "On a Diffuse Reflection of the α-Particles," Proc. R. Soc. A 82 , 495 (1909).

[7] H. Geiger, "The Scattering of the α Particles by Matter," Proc. R. Soc. A 83 , 492 (1910).

[8] E. Rutherford, "The Origin of α and β Rays From Radioactive Substances," Philos. Mag. 24 , 453 (1912).

[9] H. Geiger and E. Marsden, "The Laws of Deflexion of α Particles Through Large Angles," Philos. Mag. 25 , 604 (1913).

Rutherford Scattering

Set initial velocity of alpha:  

Discovery of the Nucleus

Rutherford discovered the nucleus by scattering alpha particles from a gold atom, expecting to find only small angle scattering based on his belief in Thomson's plum pudding model .

He was amazed to find occasional very large scattering angles, only explicable by assuming a very compact heavy central charged object: a nucleus.

For smaller nuclei, he found the alphas actually bounced off the nuclear surface, giving a different scattering pattern, and making it possible to estimate nuclear size. The nuclear radius turned out to be a fraction of a ten-thousandth of the atomic radius.

All the details can be found in my lecture on Rutherford Scattering.

Note that in the above applet, a few percent of the alphas are scattered backwards. In the real experiment, if was of order one in a thousand. This was very tedious to watch! The incoming alphas pass at random distances from the nucleus, but we've cut back to those passing relatively close. (Remember in reality, the next nucleus is around ten thousand nuclear radii away.)

30.2 Discovery of the Parts of the Atom: Electrons and Nuclei

Learning objectives.

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

• Describe how electrons were discovered.
• Explain the Millikan oil drop experiment.
• Describe Rutherford’s gold foil experiment.
• Describe Rutherford’s planetary model of the atom.

Just as atoms are a substructure of matter, electrons and nuclei are substructures of the atom. The experiments that were used to discover electrons and nuclei reveal some of the basic properties of atoms and can be readily understood using ideas such as electrostatic and magnetic force, already covered in previous chapters.

Charges and Electromagnetic Forces

In previous discussions, we have noted that positive charge is associated with nuclei and negative charge with electrons. We have also covered many aspects of the electric and magnetic forces that affect charges. We will now explore the discovery of the electron and nucleus as substructures of the atom and examine their contributions to the properties of atoms.

The Electron

Gas discharge tubes, such as that shown in Figure 30.4 , consist of an evacuated glass tube containing two metal electrodes and a rarefied gas. When a high voltage is applied to the electrodes, the gas glows. These tubes were the precursors to today’s neon lights. They were first studied seriously by Heinrich Geissler, a German inventor and glassblower, starting in the 1860s. The English scientist William Crookes, among others, continued to study what for some time were called Crookes tubes, wherein electrons are freed from atoms and molecules in the rarefied gas inside the tube and are accelerated from the cathode (negative) to the anode (positive) by the high potential. These “ cathode rays ” collide with the gas atoms and molecules and excite them, resulting in the emission of electromagnetic (EM) radiation that makes the electrons’ path visible as a ray that spreads and fades as it moves away from the cathode.

Gas discharge tubes today are most commonly called cathode-ray tubes , because the rays originate at the cathode. Crookes showed that the electrons carry momentum (they can make a small paddle wheel rotate). He also found that their normally straight path is bent by a magnet in the direction expected for a negative charge moving away from the cathode. These were the first direct indications of electrons and their charge.

The English physicist J. J. Thomson (1856–1940) improved and expanded the scope of experiments with gas discharge tubes. (See Figure 30.5 and Figure 30.6 .) He verified the negative charge of the cathode rays with both magnetic and electric fields. Additionally, he collected the rays in a metal cup and found an excess of negative charge. Thomson was also able to measure the ratio of the charge of the electron to its mass, q e q e / m e / m e —an important step to finding the actual values of both q e q e and m e m e . Figure 30.7 shows a cathode-ray tube, which produces a narrow beam of electrons that passes through charging plates connected to a high-voltage power supply. An electric field E E is produced between the charging plates, and the cathode-ray tube is placed between the poles of a magnet so that the electric field E E is perpendicular to the magnetic field B B of the magnet. These fields, being perpendicular to each other, produce opposing forces on the electrons. As discussed for mass spectrometers in More Applications of Magnetism , if the net force due to the fields vanishes, then the velocity of the charged particle is v = E / B v = E / B . In this manner, Thomson determined the velocity of the electrons and then moved the beam up and down by adjusting the electric field.

To see how the amount of deflection is used to calculate q e / m e q e / m e , note that the deflection is proportional to the electric force on the electron:

But the vertical deflection is also related to the electron’s mass, since the electron’s acceleration is

The value of F F is not known, since q e q e was not yet known. Substituting the expression for electric force into the expression for acceleration yields

Gathering terms, we have

The deflection is analyzed to get a a , and E E is determined from the applied voltage and distance between the plates; thus, q e m e q e m e can be determined. With the velocity known, another measurement of q e m e q e m e can be obtained by bending the beam of electrons with the magnetic field. Since F mag = q e vB = m e a F mag = q e vB = m e a , we have q e / m e = a / vB q e / m e = a / vB . Consistent results are obtained using magnetic deflection.

What is so important about q e / m e q e / m e , the ratio of the electron’s charge to its mass? The value obtained is

This is a huge number, as Thomson realized, and it implies that the electron has a very small mass. It was known from electroplating that about 10 8 C/kg 10 8 C/kg is needed to plate a material, a factor of about 1000 less than the charge per kilogram of electrons. Thomson went on to do the same experiment for positively charged hydrogen ions (now known to be bare protons) and found a charge per kilogram about 1000 times smaller than that for the electron, implying that the proton is about 1000 times more massive than the electron. Today, we know more precisely that

where q p q p is the charge of the proton and m p m p is its mass. This ratio (to four significant figures) is 1836 times less charge per kilogram than for the electron. Since the charges of electrons and protons are equal in magnitude, this implies m p = 1836 m e m p = 1836 m e .

Thomson performed a variety of experiments using differing gases in discharge tubes and employing other methods, such as the photoelectric effect, for freeing electrons from atoms. He always found the same properties for the electron, proving it to be an independent particle. For his work, the important pieces of which he began to publish in 1897, Thomson was awarded the 1906 Nobel Prize in Physics. In retrospect, it is difficult to appreciate how astonishing it was to find that the atom has a substructure. Thomson himself said, “It was only when I was convinced that the experiment left no escape from it that I published my belief in the existence of bodies smaller than atoms.”

Thomson attempted to measure the charge of individual electrons, but his method could determine its charge only to the order of magnitude expected.

Since Faraday’s experiments with electroplating in the 1830s, it had been known that about 100,000 C per mole was needed to plate singly ionized ions. Dividing this by the number of ions per mole (that is, by Avogadro’s number), which was approximately known, the charge per ion was calculated to be about 1 . 6 × 10 − 19 C 1 . 6 × 10 − 19 C , close to the actual value.

An American physicist, Robert Millikan (1868–1953) (see Figure 30.8 ), decided to improve upon Thomson’s experiment for measuring q e q e and was eventually forced to try another approach, which is now a classic experiment performed by students. The Millikan oil drop experiment is shown in Figure 30.9 .

In the Millikan oil drop experiment, fine drops of oil are sprayed from an atomizer. Some of these are charged by the process and can then be suspended between metal plates by a voltage between the plates. In this situation, the weight of the drop is balanced by the electric force:

The electric field is produced by the applied voltage, hence, E = V / d E = V / d , and V V is adjusted to just balance the drop’s weight. The drops can be seen as points of reflected light using a microscope, but they are too small to directly measure their size and mass. The mass of the drop is determined by observing how fast it falls when the voltage is turned off. Since air resistance is very significant for these submicroscopic drops, the more massive drops fall faster than the less massive, and sophisticated sedimentation calculations can reveal their mass. Oil is used rather than water, because it does not readily evaporate, and so mass is nearly constant. Once the mass of the drop is known, the charge of the electron is given by rearranging the previous equation:

where d d is the separation of the plates and V V is the voltage that holds the drop motionless. (The same drop can be observed for several hours to see that it really is motionless.) By 1913 Millikan had measured the charge of the electron q e q e to an accuracy of 1%, and he improved this by a factor of 10 within a few years to a value of − 1 . 60 × 10 − 19 C − 1 . 60 × 10 − 19 C . He also observed that all charges were multiples of the basic electron charge and that sudden changes could occur in which electrons were added or removed from the drops. For this very fundamental direct measurement of q e q e and for his studies of the photoelectric effect, Millikan was awarded the 1923 Nobel Prize in Physics.

With the charge of the electron known and the charge-to-mass ratio known, the electron’s mass can be calculated. It is

Substituting known values yields

where the round-off errors have been corrected. The mass of the electron has been verified in many subsequent experiments and is now known to an accuracy of better than one part in one million. It is an incredibly small mass and remains the smallest known mass of any particle that has mass. (Some particles, such as photons, are massless and cannot be brought to rest, but travel at the speed of light.) A similar calculation gives the masses of other particles, including the proton. To three digits, the mass of the proton is now known to be

which is nearly identical to the mass of a hydrogen atom. What Thomson and Millikan had done was to prove the existence of one substructure of atoms, the electron, and further to show that it had only a tiny fraction of the mass of an atom. The nucleus of an atom contains most of its mass, and the nature of the nucleus was completely unanticipated.

Another important characteristic of quantum mechanics was also beginning to emerge. All electrons are identical to one another. The charge and mass of electrons are not average values; rather, they are unique values that all electrons have. This is true of other fundamental entities at the submicroscopic level. All protons are identical to one another, and so on.

The Nucleus

Here, we examine the first direct evidence of the size and mass of the nucleus. In later chapters, we will examine many other aspects of nuclear physics, but the basic information on nuclear size and mass is so important to understanding the atom that we consider it here.

Nuclear radioactivity was discovered in 1896, and it was soon the subject of intense study by a number of the best scientists in the world. Among them was New Zealander Lord Ernest Rutherford, who made numerous fundamental discoveries and earned the title of “father of nuclear physics.” Born in Nelson, Rutherford did his postgraduate studies at the Cavendish Laboratories in England before taking up a position at McGill University in Canada where he did the work that earned him a Nobel Prize in Chemistry in 1908. In the area of atomic and nuclear physics, there is much overlap between chemistry and physics, with physics providing the fundamental enabling theories. He returned to England in later years and had six future Nobel Prize winners as students. Rutherford used nuclear radiation to directly examine the size and mass of the atomic nucleus. The experiment he devised is shown in Figure 30.10 . A radioactive source that emits alpha radiation was placed in a lead container with a hole in one side to produce a beam of alpha particles, which are a type of ionizing radiation ejected by the nuclei of a radioactive source. A thin gold foil was placed in the beam, and the scattering of the alpha particles was observed by the glow they caused when they struck a phosphor screen.

Alpha particles were known to be the doubly charged positive nuclei of helium atoms that had kinetic energies on the order of 5 MeV 5 MeV when emitted in nuclear decay, which is the disintegration of the nucleus of an unstable nuclide by the spontaneous emission of charged particles. These particles interact with matter mostly via the Coulomb force, and the manner in which they scatter from nuclei can reveal nuclear size and mass. This is analogous to observing how a bowling ball is scattered by an object you cannot see directly. Because the alpha particle’s energy is so large compared with the typical energies associated with atoms ( MeV MeV versus eV eV ), you would expect the alpha particles to simply crash through a thin foil much like a supersonic bowling ball would crash through a few dozen rows of bowling pins. Thomson had envisioned the atom to be a small sphere in which equal amounts of positive and negative charge were distributed evenly. The incident massive alpha particles would suffer only small deflections in such a model. Instead, Rutherford and his collaborators found that alpha particles occasionally were scattered to large angles, some even back in the direction from which they came! Detailed analysis using conservation of momentum and energy—particularly of the small number that came straight back—implied that gold nuclei are very small compared with the size of a gold atom, contain almost all of the atom’s mass, and are tightly bound. Since the gold nucleus is several times more massive than the alpha particle, a head-on collision would scatter the alpha particle straight back toward the source. In addition, the smaller the nucleus, the fewer alpha particles that would hit one head on.

Although the results of the experiment were published by his colleagues in 1909, it took Rutherford two years to convince himself of their meaning. Like Thomson before him, Rutherford was reluctant to accept such radical results. Nature on a small scale is so unlike our classical world that even those at the forefront of discovery are sometimes surprised. Rutherford later wrote: “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. On consideration, I realized that this scattering backwards ... [meant] ... the greatest part of the mass of the atom was concentrated in a tiny nucleus.” In 1911, Rutherford published his analysis together with a proposed model of the atom. The size of the nucleus was determined to be about 10 − 15 m 10 − 15 m , or 100,000 times smaller than the atom. This implies a huge density, on the order of 10 15 g/cm 3 10 15 g/cm 3 , vastly unlike any macroscopic matter. Also implied is the existence of previously unknown nuclear forces to counteract the huge repulsive Coulomb forces among the positive charges in the nucleus. Huge forces would also be consistent with the large energies emitted in nuclear radiation.

The small size of the nucleus also implies that the atom is mostly empty inside. In fact, in Rutherford’s experiment, most alphas went straight through the gold foil with very little scattering, since electrons have such small masses and since the atom was mostly empty with nothing for the alpha to hit. There were already hints of this at the time Rutherford performed his experiments, since energetic electrons had been observed to penetrate thin foils more easily than expected. Figure 30.11 shows a schematic of the atoms in a thin foil with circles representing the size of the atoms (about 10 − 10 m 10 − 10 m ) and dots representing the nuclei. (The dots are not to scale—if they were, you would need a microscope to see them.) Most alpha particles miss the small nuclei and are only slightly scattered by electrons. Occasionally, (about once in 8000 times in Rutherford’s experiment), an alpha hits a nucleus head-on and is scattered straight backward.

Based on the size and mass of the nucleus revealed by his experiment, as well as the mass of electrons, Rutherford proposed the planetary model of the atom . The planetary model of the atom pictures low-mass electrons orbiting a large-mass nucleus. The sizes of the electron orbits are large compared with the size of the nucleus, with mostly vacuum inside the atom. This picture is analogous to how low-mass planets in our solar system orbit the large-mass Sun at distances large compared with the size of the sun. In the atom, the attractive Coulomb force is analogous to gravitation in the planetary system. (See Figure 30.12 .) Note that a model or mental picture is needed to explain experimental results, since the atom is too small to be directly observed with visible light.

Rutherford’s planetary model of the atom was crucial to understanding the characteristics of atoms, and their interactions and energies, as we shall see in the next few sections. Also, it was an indication of how different nature is from the familiar classical world on the small, quantum mechanical scale. The discovery of a substructure to all matter in the form of atoms and molecules was now being taken a step further to reveal a substructure of atoms that was simpler than the 92 elements then known. We have continued to search for deeper substructures, such as those inside the nucleus, with some success. In later chapters, we will follow this quest in the discussion of quarks and other elementary particles, and we will look at the direction the search seems now to be heading.

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Rutherford scattering.

How did Rutherford figure out the structure of the atom without being able to see it? Simulate the famous experiment in which he disproved the Plum Pudding model of the atom by observing alpha particles bouncing off atoms and determining that they must have a small core.

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Bridge Course Class 9th Science

Course: bridge course class 9th science   >   unit 2, rutherford scattering experiment.

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Video transcript

• Structure of Atom
• Rutherford Atomic Model And Its Limitations

Rutherford Atomic Model and Limitations

Define rutherford atomic model.

Rutherford Atomic Model – The plum pudding model given by J. J. Thomson failed to explain certain experimental results associated with the atomic structure of elements. Ernest Rutherford, a British scientist conducted an experiment and based on the observations of this experiment, he explained the atomic structure of elements and proposed Rutherford’s Atomic Model.

Table of Contents

• Rutherfords Alpha Scattering Experiment

Observations of Rutherford’s Alpha Scattering Experiment

Rutherford atomic model, limitations of rutherford atomic model, recommended videos, frequently asked questions – faqs.

Rutherford’s Alpha Scattering Experiment

Rutherford conducted an experiment by bombarding a thin sheet of gold with α-particles and then studied the trajectory of these particles after their interaction with the gold foil.

Rutherford, in his experiment, directed high energy streams of α-particles from a radioactive source at a thin sheet (100 nm thickness) of gold. In order to study the deflection caused to the α-particles, he placed a fluorescent zinc sulphide screen around the thin gold foil. Rutherford made certain observations that contradicted Thomson’s atomic model .

The observations made by Rutherford led him to conclude that:

• A major fraction of the α-particles bombarded towards the gold sheet passed through the sheet without any deflection, and hence most of the space in an atom is empty .
• Some of the α-particles were deflected by the gold sheet by very small angles, and hence the positive charge in an atom is not uniformly distributed . The positive charge in an atom is concentrated in a very small volume .
• Very few of the α-particles were deflected back, that is only a few α-particles had nearly 180 o angle of deflection. So the volume occupied by the positively charged particles in an atom is very small as compared to the total volume of an atom .

Based on the above observations and conclusions, Rutherford proposed the atomic structure of elements. According to the Rutherford atomic model:

• The positive charge and most of the mass of an atom is concentrated in an extremely small volume. He called this region of the atom as a nucleus.
• Rutherford’s model proposed that the negatively charged electrons surround the nucleus of an atom. He also claimed that the electrons surrounding the nucleus revolve around it with very high speed in circular paths. He named these circular paths as orbits.
• Electrons being negatively charged and nucleus being a densely concentrated mass of positively charged particles are held together by a strong electrostatic force of attraction.

Although the Rutherford atomic model was based on experimental observations, it failed to explain certain things.

• Rutherford proposed that the electrons revolve around the nucleus in fixed paths called orbits. According to Maxwell, accelerated charged particles emit electromagnetic radiations and hence an electron revolving around the nucleus should emit electromagnetic radiation. This radiation would carry energy from the motion of the electron which would come at the cost of shrinking of orbits. Ultimately the electrons would collapse in the nucleus. Calculations have shown that as per the Rutherford model, an electron would collapse into the nucleus in less than 10 -8 seconds. So the Rutherford model was not in accordance with Maxwell’s theory and could not explain the stability of an atom .
• One of the drawbacks of the Rutherford model was also that he did not say anything about the arrangement of electrons in an atom which made his theory incomplete.
• Although the early atomic models were inaccurate and failed to explain certain experimental results, they formed the base  for future developments in the world of quantum mechanics .

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The Gold Foil Experiment

Structure of Atom Class 11 Chemistry

Drawbacks of Rutherford Atomic Model

What was the speciality of Rutherford’s atomic model?

Rutherford was the first to determine the presence of a nucleus in an atom. He bombarded α-particles on a gold sheet, which made him encounter the presence of positively charged specie inside the atom.

What is Rutherford’s atomic model?

Rutherford proposed the atomic structure of elements. He explained that a positively charged particle is present inside the atom, and most of the mass of an atom is concentrated over there. He also stated that negatively charged particles rotate around the nucleus, and there is an electrostatic force of attraction between them.

What are the limitations of Rutherford’s atomic model?

Rutherford failed to explain the arrangement of electrons in an atom. Like Maxwell, he was unable to explain the stability of the atom.

What kind of experiment did Rutherford’s perform?

Rutherford performed an alpha scattering experiment. He bombarded α-particles on a gold sheet and then studied the trajectory of these α-particles.

What was the primary observation of Rutherford’s atomic model?

Rutherford observed that a microscopic positively charged particle is present inside the atom, and most of the mass of an atom is concentrated over there.

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I am very happy with the answer that I obtained, however Ernest Rutherford’s Atomic Model never had any neutrons in the nucleus. James Chadwick discovered the neutron later in 1932. However, the limitations and observations of his theory on this web page seem to be correct.

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Rutherford’s Alpha Scattering Experiment

Rutherford’s Alpha Scattering Experiment is the fundamental experiment done by Earnest Rutherford’s Alpha Scattering Experiment that gives the fundamental about the structure of the atom. Rutherford in his experiment directed high-energy streams of α-particles from a radioactive source at a thin sheet (100 nm thickness) of gold. Then the deflection of these alpha particles tells us about the structure of atoms.

In this article, we will study about constituents of atoms, Rutherford’s  Alpha Scattering Experiment,

What are Constituents of an Atom?

An atom consists of Electrons, Protons, and Neutrons are the fundamental particles or sub-atomic particles that build the structure of an atom. Let us understand each term.

• Electron: In 1897, J. J. Thomson discovered negatively charged particles towards the anode, these rays are emitted by the cathode in a cathode ray experiment. Then these negatively charged particles are proposed as Electrons .
• Protons: In 1886, Ernest Goldstein discovered that an anode emitted positively charged particles with a different condition in the same tube,  known as Canal rays or as Protons .
• Neutrons: A subatomic particle with no charge and a mass equivalent to protons in the nucleus of all atoms was discovered by J. Chadwick. These neutrally charged particles are termed Neutrons .

The image added below shows the structure of an atom.

Learn more about, Atomic Structure

Structure of Atom

Isotopes are the elements that have the same atomic number but different mass. e.g. Isotopes of the Hydrogen atoms are Protium ( 1 H 1 ), Deuterium ( 2 H 1 ) and Tritium( 3 H 1 ). Isotopes of the Carbon atoms are 12 C 6 , 13 C 6 , 14 C 6 .

Isobars are the elements that have different atomic number but have same mass number. e.g. 19 K 40 , 18 Ar 40 , 20 Ca 40 , here all the elements having same mass number hence they are isobars.

He conduct an experiment by bombarding alpha particles into a thin sheet of gold and then notices their interaction with the gold foil and trajectory or path followed by these particles.

In the experiment, Rutherford passes very high streams of alpha-particles from a radioactive source i.e. alpha-particle emitter, at a thin sheet of100 nm thickness of gold. In order to examine the deflection produced by the alpha particles, he placed a screen of fluorescent zinc sulphide around the thin gold foil. Rutherford made certain observations that oppose Thomson’s atomic model.

Observations of Rutherford’s Alpha Scattering Experiment

The observations of Rutherford’s Alpha Scattering Experiment are:

• First, he observe that most of the α-particles that are bombarded towards the gold sheet pass away the foil without any deflection, and hence it shows most of the space is empty.
• Out of all, some of the α-particles were deflected through the gold sheet by very small angles, and hence it shows the positive charge in an atom is non-uniformly distributed. The positive charge is concentrated in a very small volume in an atom.
• Very few of the alpha-particles(1-2%) were deflected back, i.e. only a very less amount of α-particles had nearly 180° angle of deflection. this shows that the volume occupied by the positively charged particles is very small as compared to the total volume of an atom.

Rutherford Atomic Model

Rutherford proposed the atomic structure of elements, on the basis of his experiment. According to Rutherford’s atomic model:

• Positively charged particle was concentrated in an extremely small volume and most of the mass of an atom was also in that volume. He called this a nucleus of an atom.
• Rutherford proposed that there is negatively charged electrons around the nucleus of an atom. the electron surrounding the nucleus revolves around it in a circular path with very high speed. He named orbits to these circular paths.
• Nucleus being a densely concentrated mass of positively charged particles and electrons being negatively charged are held together by a strong force of attraction called electrostatic forces of attraction.

Learn about, Rutherford Atomic Model

Limitations of Rutherford Atomic Model

The Rutherford atomic model is failed to explain certain things.

• According to Maxwell, an electron revolving around the nucleus should emit electromagnetic radiation due to accelerated charged particles emit electromagnetic radiation. but Rutherford model says that the electrons revolve around the nucleus in fixed paths called orbits. The radiation would carry energy from the motion which led to the shrinking of orbit. Ultimately electrons would collapse inside the nucleus.
• As per the Rutherford model, calculations have shown that an electron would collapse in the nucleus in less than 10 -8 seconds. So Rutherford model has created a high contradiction with Maxwell’s theory and Rutherford later could not explain the stability of an atom.
• Rutherford also did not describe the arrangement of electrons in the orbit as one of the other drawbacks of his model.

Regardless of seeing the early atomic models were inaccurate and failed to explain certain experimental results, they were the base for future developments in the world of quantum mechanics.

Sample Questions on Rutherford’s Alpha Scattering Experiment

Some sample questions on Rutherford’s Alpha Scattering Experiment is,

Q1: Represent Element ‘X’ which contains 15 electrons and 16 neutrons.

Atomic number of element = No. of electron = 15 Mass number of element = no. of electrons + no. of neutrons = 15 + 16 = 31 Correct representation of element X is 31 X 15 .

Q2: Name particle and give its location in the atom which has no charge and has a mass nearly equal to that of a proton.

The particle which has no charge and has a mass nearly equal to that of a proton is a neutron and it is present in the nucleus of the atom.

Q3: An atom has both electron attribute negative charge and protons attribute positive charge but why there is no charge?

Positive and negative charges of protons and electrons are equal in magnitude, they cancel the effect of each other. So, the atom as a whole is electrically neutral.

Q4: What is Valency of Sodium Atom (Na)?

The atomic number of sodium = 11. Electronic configuration (2, 8, 1). By losing one electron it gains stability hence its valency is 1.

Q5: Which property do the following pairs show? 209 X 84 and 210 X 84

Atomic number of X is the same hence the pair shows an isotopic property. So, 209 X 84 and 210 X 84 are isotopes.

Dalton’s Atomic Theory Thomson’s Atomic Model Quantum Numbers

Rutherford’s Alpha Scattering Experiment FAQs

What is name of atom which has one electron, one proton and no neutron.

Atom with one electron, one proton and no neutron is Hydrogen, ( 1 H 1 ).

What is Ground State of an Atom?

It is the state of an atom where all the electrons in the atom are in their lowest energy state or levels is called the ground state.

What was Rutherford’s Alpha Particle Scattering Experiment?

Rutherford’s Alpha Particle Scattering Experiment is the fundamental experiment that gives the basic structure of an atom.

What was Conclusion of Rutherford’s Alpha Scattering Experiment?

Conclusion of Rutherford’s Alpha Scattering Experiment is, Atom is largely empty and has a heavy positive-charged body at the center called the nucleus. The central nucleus is positively charged and the negatively-charged electrons revolve around the nucleus.

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Rutherford Scattering ( AQA GCSE Physics )

Revision note.

Physics Project Lead

Rutherford Scattering

Alpha scattering.

• Physicist, Ernest Rutherford was instructing two of his students, Hans Geiger and Ernest Marsden to carry out the experiment
• They were directing a beam of alpha particles (He 2+ ions) at a thin gold foil
• They expected the alpha particles to travel through the gold foil, and maybe change direction a small amount
• Most of the alpha particles passed straight through the foil
• Some of the alpha particles changed direction but continued through the foil
• A few of the alpha particles bounced back off the gold foil
• The bouncing back could not be explained by the Plum Pudding model, so a new model had to be created

When alpha particles are fired at thin gold foil, most of them go straight through, some are deflected and a very small number bounce straight back

The Nuclear Model

• Ernest Rutherford made different conclusions from the findings of the experiment
• The table below describes the findings and conclusions of A, B and C from the image above:

Alpha Scattering Findings and Conclusions Table

• Rutherford proposed the nuclear model of the atom
• Nearly all of the mass of the atom is concentrated in the centre of the atom (in the nucleus)
• The nucleus is positively charged
• Negatively charged electrons orbit the nucleus at a distance
• The nuclear model could explain experimental observations better than the Plum Pudding model

The Nuclear model replaced the Plum Pudding model as it could better explain the observations of Rutherford’s Scattering Experiment

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