DISCOVERY OF ATOM

by the beginning of the 19th century, an important fact was well established—the masses of reactants in specific chemical reactions always have a particular mass ratio. This is very strong indirect evidence that there are basic units (atoms and molecules) that have these same mass ratios. The English chemist John Dalton (1766–1844) did much of this work, with significant contributions by the Italian physicist Amedeo Avogadro (1776–1856). It was Avogadro who developed the idea of a fixed number of atoms and molecules in a mole, and this special number is called Avogadro’s number in his honor.

Law of Constant Composition

Joseph Proust (1754-1826) formulated the law of constant composition (also called the law of definite proportions). This law states that if a compound is broken down into its constituent elements, the masses of the constituents will always have the same proportions, regardless of the quantity or source of the original substance. Joseph Proust based this law primarily on his experiments with basic copper carbonate. The illustration below depicts this law; 31 grams of H2O and 8 grams of H2O are made up of the same percent of hydrogen and oxygen.

Law of multiple proportions, statement that when two elements combine with each other to form more than one compound, the weights of one element that combine with a fixed weight of the other are in a ratio of small whole numbers. For example, there are five distinct oxides of nitrogen, and the weights of oxygen in combination with 14 grams of nitrogen are, in increasing order, 8, 16, 24, 32, and 40 grams, or in a ratio of 1, 2, 3, 4, 5. The law was announced (1803) by the English chemist John Dalton, and its confirmation for a wide range of compounds served as the most powerful argument in support of Dalton’s theory that matter consists of indivisible atoms.

Law of multiple proportions, statement that when two elements combine with each other to form more than one compound, the weights of one element that combine with a fixed weight of the other are in a ratio of small whole numbers. For example, there are five distinct oxides of nitrogen, and the weights of oxygen in combination with 14 grams of nitrogen are, in increasing order, 8, 16, 24, 32, and 40 grams, or in a ratio of 1, 2, 3, 4, 5. The law was announced (1803) by the English chemist John Dalton, and its confirmation for a wide range of compounds served as the most powerful argument in support of Dalton’s theory that matter consists of indivisible atoms.

Dalton's atomic theory

Key Points

• 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.

Dalton's atomic theory
Part 1: All matter is made of atoms.
Dalton hypothesized that the law of conservation of mass and the law of definite proportions could be explained using the idea of atoms. He proposed that all matter is made of tiny indivisible particles called atoms, which he imagined as "solid, massy, hard, impenetrable, movable particle(s)".
It is important to note that since Dalton did not have the necessary instruments to see or otherwise experiment on individual atoms, he did not have any insight into whether they might have any internal structure. We might visualize Dalton's atom as a piece in a molecular modeling kit, where different elements are spheres of different sizes and colors. While this is a handy model for some applications, we now know that atoms are far from being solid spheres.
Part 2: All atoms of a given element are identical in mass and properties.
Dalton proposed that every single atom of an element, such as gold, is the same as every other atom of that element. He also noted that the atoms of one element differ from the atoms of all other elements. Today, we still know this to be mostly true. A sodium atom is different from a carbon atom. Elements may share some similar boiling points, melting points, and electronegativities, but no two elements have the same exact set of properties.

• Part 3: Compounds are combinations of two or more different types of atoms.
  In the third part of Dalton's atomic theory, he proposed that compounds are combinations of two or more different types of atoms. An example of such a compound is table salt. Table salt is a combination of two separate elements with unique physical and chemical properties. The first, sodium, is a highly reactive metal. The second, chlorine, is a toxic gas. When they react, the atoms combine in a 1:1 ratio to form white crystals of NaCl , which we can sprinkle on our food .

Since atoms are indivisible, they will always combine in simple whole number ratios. Therefore, it would not make sense to write a formula such as Na 0.5 Cl 0.5 because you can't have half of an atom!

• Part 4: A chemical reaction is a rearrangement of atoms.
  In the fourth and final part of Dalton's atomic theory, he suggested that chemical reactions don't destroy or create atoms. They merely rearranged the atoms. Using our salt example again, when sodium combines with chlorine to make salt, both the sodium and chlorine atoms still exist. They simply rearrange to form a new compound.

Discovering Electrons
The first cathode-ray tube (CRT) was invented by Michael Faraday (1791-1867). Cathode rays are a type of radiation emitted by the negative terminal, the cathode, and were discovered by passing electricity through nearly-evacuated glass tubes. The radiation crosses the evacuated tube to the positive terminal, the anode. Cathode rays produced by the CRT are invisible and can only be detected by light emitted by the materials that they strike, called phosphors, painted at the end of the CRT to reveal the path of the cathode rays. These phosphors showed that cathode rays travel in straight lines and have properties independent of the cathode material (whether it is gold, silver, etc.). Another significant property of cathode rays is that they are deflected by magnetic and electric fields in a manner that is identical to negatively charged material. Due to these observations, J.J. Thompson (1856-1940) concluded that cathode rays are negatively charged particles that are located in all atoms. It was George Stoney who first gave the term electrons to the cathode rays. The below figures depict the way that the cathode ray is effected by magnetics. The cathode ray is always attracted by the positive magnet and deflected by the negative magnets.

To account for these observations, Rutherford devised a model called the nuclear atom. In this model, the positive charge is held in an extremely small area called the nucleus, located in the middle of the atom. Outside of the nucleus the atom is largely composed of empty space. This model states that there were positive particles within the nucleus, but failed to define what these particles are. Rutherford discovered these particles in 1919, when he conducted an experiment that scattered alpha particles against nitrogen atoms. When the alpha particles and nitrogen atoms collided protons were released.

The Discovery of the Neutron

In 1933, James Chadwick (1891-1974) discovered a new type of radiation that consisted of neutral particles. It was discovered that these neutral atoms come from the nucleus of the atom. This last discovery completed the atomic model.

Atomic nucleus

The atomic nucleus is the small, dense region consisting of protons and neutrons at the center of an atom, discovered in 1911 by Ernest Rutherford.An atom is composed of a positively-charged nucleus, with a cloud of negatively-charged electrons surrounding it, bound together by electrostatic force. Almost all of the mass of an atom is located in the nucleus, with a very small contribution from the electron cloud. Protons and neutrons are bound together to form a nucleus by the nuclear force.

History

Ernest Rutherford later devised an experiment with his research partner Hans Geiger and with help of Ernest Marsden, that involved the deflection of alpha particles (helium nuclei) directed at a thin sheet of metal foil. He reasoned that if J.J Thomson's model were correct, the positively charged alpha particles would easily pass through the foil with very little deviation in their paths, as the foil should act as electrically neutral if the negative and positive charges are so intimately mixed as to make it appear neutral. To his surprise, many of the particles were deflected at very large angles. Because the mass of an alpha particle is about 8000 times that of an electron, it became apparent that a very strong force must be present if it could deflect the massive and fast moving alpha particles. He realized that the plum pudding model could not be accurate and that the deflections of the alpha particles could only be explained if the positive and negative charges were separated from each other and that the mass of the atom was a concentrated point of positive charge. This justified the idea of a nuclear atom with a dense center of positive charge and mass.

Forces
Nuclei are bound together by the residual strong force (nuclear force). The residual strong force is a minor residuum of the strong interaction which binds quarks together to form protons and neutrons. This force is much weaker between neutrons and protons because it is mostly neutralized within them, in the same way that electromagnetic forces between neutral atoms (such as van der Waals forces that act between two inert gas atoms) are much weaker than the electromagnetic forces that hold the parts of the atoms together internally (for example, the forces that hold the electrons in an inert gas atom bound to its nucleus).
The nuclear force is highly attractive at the distance of typical nucleon separation, and this overwhelms the repulsion between protons due to the electromagnetic force, thus allowing nuclei to exist. However, the residual strong force has a limited range because it decays quickly with distance (see Yukawa potential); thus only nuclei smaller than a certain size can be completely stable. The largest known completely stable nucleus (i.e. stable to alpha, beta, and gamma decay) is lead-208 which contains a total of 208 nucleons (126 neutrons and 82 protons). Nuclei larger than this maximum are unstable and tend to be increasingly short-lived with larger numbers of nucleons. However, bismuth-209 is also stable to beta decay and has the longest half-life to alpha decay of any known isotope, estimated at a billion times longer than the age of the universe.
The residual strong force is effective over a very short range (usually only a few femtometres (fm); roughly one or two nucleon diameters) and causes an attraction between any pair of nucleons. For example, between protons and neutrons to form [NP] deuteron, and also between protons and protons, and neutrons and neutrons.

Atomic number, the number of a chemical element in the periodic system, whereby the elements are arranged in order of increasing number of protons in the nucleus. Accordingly, the number of protons, which is always equal to the number of electrons in the neutral atom, is also the atomic number. An atom of iron has 26 protons in its nucleus; therefore the atomic number of iron is 26.In the symbol representing a particular nuclear or atomic species, the atomic number may be indicated as a left subscript. An atom or a nucleus of iron (chemical symbol Fe), for example, may be written ₂₆Fe.

Mass number, is, the sum of the numbers of protons and neutrons present in the nucleus of an atom. The mass number is commonly cited in distinguishing among the isotopes of an element, all of which have the same atomic number (number of protons) and are represented by the same literal symbol; for example, the two best known isotopes of uranium (those with mass numbers 235 and 238) are designated uranium-235 (symbolized ²³⁵U) and uranium-238 (²³⁸U).: mass number = protons + neutrons.

isotope, one of two or more species of atoms of a chemical element with the same atomic number and position in the periodic table and nearly identical chemical behaviour but with different atomic masses and physical properties. Every chemical element has one or more isotopes.
An atom is first identified and labeled according to the number of protons in its nucleus. This atomic number is ordinarily given the symbol Z. The great importance of the atomic number derives from the observation that all atoms with the same atomic number have nearly, if not precisely, identical chemical properties. A large collection of atoms with the same atomic number constitutes a sample of an element. A bar of pure uranium, for instance, would consist entirely of atoms with atomic number 92. The periodic table of the elements assigns one place to every atomic number, and each of these places is labeled with the common name of the element, as, for example, calcium, radon, or uranium.Not all the atoms of an element need have the same number of neutrons in their nuclei. In fact, it is precisely the variation in the number of neutrons in the nuclei of atoms that gives rise to isotopes. Hydrogen is a case in point. It has the atomic number 1. Three nuclei with one proton are known that contain 0, 1, and 2 neutrons, respectively. The three share the place in the periodic table assigned to atomic number 1 and hence are called isotopes (from the Greek isos, meaning “same,” and topos, signifying “place”) of hydrogen.Many important properties of an isotope depend on its mass. The total number of neutrons and protons (symbol A), or mass number, of the nucleus gives approximately the mass measured on the so-called atomic-mass-unit (amu) scale. The numerical difference between the actual measured mass of an isotope and A is called either the mass excess or the mass defect.

Example of isotope: The common examples are the isotopes of hydrogen and carbon. If we talk about the element Hydrogen, it has three stable isotopes namely protium, deuterium, and tritium. These isotopes have the same number of protons but a different number of neutrons wherein protium has zero, deuterium has one and tritium has two.

When we look at carbon it also has three isotopes namely Carbon-12, Carbon-13, and Carbon-14. The numbers 12, 13, and 14 are the isotopes’ atomic masses. Here, Carbon-12 is a stable isotope whereas carbon-14 is mostly a radioactive isotope.

Apart from these some other common isotope examples include – Tin has 22 isotopes, Zinc has 21 known isotopes, Neon is a mix of 3 isotopes, natural xenon consists of a mixture of 9 stable isotopes, Nickel has 14 known isotopes.