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    1. 1. CHEMISTRY PROJECT WORK Atoms And Molecules • Submitted by – Shivam • To – Mrs. Samnol • Class - IX ‘A’ 1
    2. 2. Chapter: Atoms and Molecules 2
    3. 3. Introduction The molecular structure hypothesis - that a molecule is a collection of atoms linked by a network of bonds - was forged in the crucible of nineteenth century experimental chemistry. It has continued to serve as the principal means of ordering and classifying the observations of chemistry. The difficulty with this hypothesis was that it was not related directly to quantum mechanics, the physics which governs the motions of the nuclei and electrons that make up the atoms and the bonds. Indeed there was, and with some there still is, a prevailing opinion that these fundamental concepts, while unquestionably useful, were beyond theoretical definition. We have in chemistry an understanding based on a classification scheme that is both powerful and at the same time, because of its empirical nature, limited. 3
    4. 4. SYMBOLS AND FORMULAS • A unique symbol is used to represent each element. • Formulas are used to represent compounds. ELEMENTAL SYMBOLS • A symbol is assigned to each element. The symbol is based on the name of the element and consists of one capital letter or a capital letter followed by a lower case letter. • Some symbols are based on the Latin or German name of the element. 4
    5. 5. 5
    6. 6. chlorine bromine iodine 6
    7. 7. COMPOUND FORMULAS • A compound formula consists of the symbols of the elements found in the compound. Each elemental symbol represents one atom of the element. If more than one atom is represented, a subscript following the elemental symbol is used. 7
    8. 8. EXAMPLES OF COMPOUND FORMULAS • Carbon monoxide, CO (one atom of C and one atom of O are represented). • Water, H2O (two atoms of H and one atom of O are represented). • Ammonia, NH3 (one atom of N and 3 atoms of H are represented). 8
    9. 9. THE STRUCTURE OF ATOMS • Atoms are made up of three subatomic particles, protons, neutrons, and electrons. • The protons and neutrons are tightly bound together to form the central portion of an atom called the nucleus. • The electrons are located outside of the nucleus and thought to move very rapidly throughout a relatively large volume of space surrounding the small but very heavy nucleus. 9
    10. 10. PROTONS • Protons are located in the nucleus of an atom. They carry a +1 electrical charge and have a mass of 1 atomic mass unit (u). NEUTRONS • Neutrons are located in the nucleus of an atom. They carry no electrical charge and have a mass of 1 atomic mass unit (u). ELECTRONS • Electrons are located outside the nucleus of an atom. They carry a -1 electrical charge and have a mass of 1/1836 atomic mass unit (u). They move rapidly around the heavy nucleus. 10
    12. 12. ATOMIC NUMBER OF AN ATOM • The atomic number of an atom is equal to the number of protons in the nucleus of the atom. • Atomic numbers are represented by the symbol Z. MASS NUMBER OF AN ATOM • The mass number of an atom is equal to the sum of the number of protons and neutrons in the nucleus of the atom. • Mass numbers are represented by the symbol A. 12
    13. 13. ISOTOPES • Isotopes are atoms that have the same number of protons in the nucleus but different numbers of neutrons. That is, they have the same atomic number but different mass numbers. • Because they have the same number of protons in the nucleus, all isotopes of the same element have the same number of electrons outside the nucleus. 13
    14. 14. SYMBOLS FOR ISOTOPES • Isotopes are represented by the symbol , where Z is the atomic number, A is the mass number and E is the elemental symbol. 60 • An example of an isotope symbol is 28 Ni . This symbol represents an isotope of nickel that contains 28 protons and 32 neutrons in the nucleus. • Isotopes are also represented by the notation: Name-A, where Name is the name of the element and A is the mass number of the isotope. • An example of this isotope notation is magnesium-26. This represents an isotope of magnesium that has a mass number of 26. Click here to play Coached Problem 14
    15. 15. RELATIVE MASSES • The extremely small size of atoms and molecules makes it inconvenient to use their actual masses for measurements or calculations. Relative masses are used instead. • Relative masses are comparisons of actual masses to each other. For example if an object had twice the mass of another object, their relative masses would be 2 to 1. 15
    16. 16. ATOMIC MASS UNIT (u) • An atomic mass unit is a unit used to express the relative masses of atoms. One atomic mass unit is equal to 1/12 the mass of a carbon-12 atom. • A carbon-12 atom has a relative mass of 12 u. • An atom with a mass equal to 1/12 the mass of a carbon-12 atom would have a relative mass of 1 u. • An atom with a mass equal to twice the mass of a carbon-12 atom would have a relative mass of 24 u. 16
    17. 17. ATOMIC WEIGHT • The atomic weight of an element is the relative mass of an average atom of the element expressed in atomic mass units. • Atomic weights are the numbers given at the bottom of the box containing the symbol of each element in the periodic table. • According to the periodic table, the atomic weight of nitrogen atoms (N) is 14.0 u, and that of silicon atoms (Si) is 28.1 u. This means that silicon atoms are very close to twice as massive as nitrogen atoms. Put another way, it means that two nitrogen atoms have a total mass very close to the mass of a single silicon atom. 17
    18. 18. MOLECULAR WEIGHT • The relative mass of a molecule in atomic mass units is called the molecular weight of the molecule. • Because molecules are made up of atoms, the molecular weight of a molecule is obtained by adding together the atomic weights of all the atoms in the molecule. • The formula for a molecule of water is H2O. This means one molecule of water contains two atoms of hydrogen, H, and one atom of oxygen, O. The molecular weight of water is then the sum of two atomic weights of H and one atomic weight of O: MW = 2(at. wt. H) + 1(at. wt. O) MW = 2(1.01 u) + 1(16.0 u) = 18.02 u 18
    19. 19. ISOTOPES AND ATOMIC WEIGHTS • Many elements occur naturally as a mixture of several isotopes. • The atomic weight of elements that occur as mixtures of isotopes is the average mass of the atoms in the isotope mixture. • The average mass of a group of atoms is obtained by dividing the total mass of the group by the number of atoms in the group. • A practical way of determining the average mass of a group of isotopes is to assume the group consists of 100 atoms and use the percentage of each isotope to represent the number of atoms of each isotope present in the group. 19
    20. 20. • The use of percentages and the mass of each isotope leads to the following equation for calculating atomic weights of elements that occur naturally as a mixture of isotopes. According to this equation, the atomic weight of an element is calculated by multiplying the percentage of each isotope in the element by the mass of the isotope, then adding the resulting products together and dividing the resulting mass by 100. 20
    21. 21. • A specific example of the use of the equation is shown below for the element boron that consists of 19.78% boron-10 with a mass of 10.01 u, and 80.22% boron-11 with a mass of 11.01u. This calculated value is seen to agree with the value given in the periodic table. Click here to play Coached Problem 21
    22. 22. THE MOLE CONCEPT APPLIED TO ELEMENTS • The number of atoms in one mole of any element is called Avogadro's number and is equal to 6.022x1023 . • A one-mole sample of any element will contain the same number of atoms as a one-mole sample of any other element. • One mole of any element is a sample of the element with a mass in grams that is numerically equal to the atomic weight of the element. EXAMPLES OF THE MOLE CONCEPT • 1 mole Na = 22.99 g Na = 6.022x1023 Na atoms • 1 mole Ca = 40.08 g Ca = 6.022x1023 Ca atoms • 1 mole S = 32.06 g S = 6.022x1023 S atoms Click here to play Chemistry Interactive 22
    23. 23. THE MOLE CONCEPT APPLIED TO COMPOUNDS • The number of molecules in one mole of any compound is called Avogadro's number and is numerically equal to 6.022x1023. • A one-mole sample of any compound will contain the same number of molecules as a one-mole sample of any other compound. • One mole of any compound is a sample of the compound with a mass in grams equal to the molecular weight of the compound. Click here to play Coached Problem 23
    24. 24. EXAMPLES OF THE MOLE CONCEPT • 1 mole H2O = 18.02 g H2O = 6.022x1023 H2O molecules • 1 mole CO2 = 44.01 g CO2 = 6.022x1023 CO2 molecules • 1 mole NH3 = 17.03 g NH3 = 6.022x1023 NH3 molecules THE MOLE AND CHEMICAL CALCULATIONS • The mole concept can be used to obtain factors that are useful in chemical calculations involving both elements and compounds. 24
    25. 25. CALCULATIONS INVOLVING ELEMENTS • The mole-based relationships given earlier as examples for elements provide factors for solving problems. • The relationships given earlier for calcium are: 1 mole Ca= 40.08 g Ca = 6.022x1023 Ca atoms. • Any two of these quantities can be used to provide factors for use in solving numerical problems. • Examples of two of the six possible factors are: and 25
    26. 26. EXAMPLE PROBLEM • Calculate the number of moles of Ca contained in a 15.84 g sample of Ca. • The solution to the problem is: • We see in the solution that the g Ca units in the denominator of the factor cancel the g Ca units in the given quantity, leaving the correct units of mole Ca for the answer. 26
    27. 27. CALCULATIONS INVOLVING COMPOUNDS • The mole concept applied earlier to molecules can be applied to the individual atoms that are contained in the molecules. • An example of this for the compound CO2 is: 1 mole CO2 molecules = 1 mole C atoms + 2 moles O atoms 44.01 g CO2 = 12.01 g C + 32.00 g O 6.022x1023 CO2 molecules = 6.022x1023 C atoms + (2) 6.022x1023 O atoms. • Any two of these nine quantities can be used to provide factors for use in solving numerical problems. 27
    28. 28. • Example 1: How many moles of O atoms are contained in 11.57 g of CO2? • Note that the factor used was obtained from two of the nine quantities given on the previous slide. 28
    29. 29. • Example 2: How many CO2 molecules are needed to contain 50.00 g of C? • Note that the factor used was obtained from two of the nine quantities given on a previous slide. 29
    30. 30. • Example 3: What is the mass percentage of C in CO 2? • The mass percentage is calculated using the following equation: • If a sample consisting of 1 mole of CO2 is used, the molebased relationships given earlier show that 1 mole CO2=44.01 g CO2=12.01 g C + 32.00 g O. Thus, the mass of C in a specific mass of CO2 is known, and the problem is solved as follows: 30
    31. 31. • Example 4: What is the mass percentage of oxygen in CO2? • The mass percentage is calculated using the following equation: • Once again, a sample consisting of 1 mole of CO2 is used to take advantage of the mole-based relationships given earlier where: 1 mole CO2 = 44.01g CO2 = 12.01 g C + 32.00g O 31
    32. 32. • Thus, the mass of O in a specific mass of CO2 is known, and the problem is solved as follows: • We see that the % C + % O = 100% , which should be the case because C and O are the only elements present in CO2. Click here to play Coached Problem 32
    33. 33. 33
    34. 34. The theory recovers the central operational concepts of the molecular structure hypothesis, that of a functional grouping of atoms with an additive and characteristic set of properties, together with a definition of the bonds that link the atoms and impart the structure. Not only does the theory thereby quantify and provide the physical understanding of the existing concepts of chemistry, it makes possible new applications of theory. These new applications will eventually enable one to perform on a computer, in a manner directly paralleling experiment, everything that can now be done in the laboratory, but more quickly and more efficiently, by linking together the functional groups of theory. These applications include the design and synthesis of new molecules and new materials with specific desirable properties. The theory of atoms in molecules enables one to take advantage of the single most important observation of chemistry, that of a functional group with a characteristic set of properties. This document outlines and illustrates the topological basis of the theory and its relation to the quantum mechanics of an open system. 34
    35. 35. • • What is an Atom? Matter is composed of atoms. This is a consequence of the manner in which the electrons are distributed throughout space in the attractive field exerted by the nuclei. The nuclei act as point attractors immersed in a cloud of negative charge, the electron density (r). The electron density describes the manner in which the electronic charge is distributed throughout real space. The electron density is a measurable property and it determines the appearance and form of matter. This is illustrated in the following figures. Figure 1 displays the spatial distribution of the electron density in the plane containing the two carbon and four hydrogen nuclei of the ethane molecule. The electron density is a maximum at the position of each nucleus and decays rapidly away from these positions. When this diagram is translated into three dimensions, the cloud of negative charge is seen to be most dense at nuclear positions and to become more diffuse as one moves away from these centers of attraction, as illustrated in Figure 2. The presence of local maxima at the positions of the nuclei is the general and also the dominant topological property of (r). Figure 3 illustrates the same feature for the 110 plane of carbon nuclei in the diamond lattice. 35
    36. 36. • To determine what physics has to say about this property of the electron density one must consider not the density itself but the field one obtains by following the trajectories traced out by the gradient vectors of the density. Starting at any point, one determines the gradient of (r). This is a vector that points in the direction of maximum increase in the density. One makes an infinitesimal step in this direction and then recalculates the gradient to obtain the new direction. By continued repetition of this process, one traces out a trajectory of (r). A gradient vector map generated in this manner is illustrated in the upper diagram of Figure 4 for the same plane of the ethane molecule shown in Figure 1. Since the density exhibits a maximum at the position of each nucleus, sets of trajectories terminate at each nucleus. The nuclei are the attractors of the gradient vector field of the electron density. Because of this fundamental property, the space of the molecule is disjoint and exhaustively partitioned into basins, a basin being the region of space traversed by the trajectories terminating at a given nucleus or attractor. Since a single attractor is associated with each basin, an atom is defined as the union of an attractor and its basin. 36
    37. 37. ATOMS AND MOLECULES • One of the most powerful developments in the history of science is the atomic/molecular model of matter, which can be used to explain and predict a large variety of phenomena. Ideas about elements, atoms, and their combination in molecules or large arrays develop from two notions: that matter is made of invisibly tiny pieces and that the enormous variety of materials in the world is the result of different combinations of a relatively small number of basic ingredients. • In high school, these ideas extend to benchmarks about the relation of a material's properties to its atomic or molecular make-up, the structure of the atom itself, and the existence of isotopes and radioactivity. The sections Understanding Fire and Splitting the Atom in Science for All Americans and Benchmarks Chapter 10: HISTORICAL PERSPECTIVES could provide context in instruction for the study of atomic/molecular theory. 37
    38. 38. • The atoms in molecules or atoms-in-molecules or quantum theory of atoms in molecules (Qtaim) approach is a quantum chemical model that characterizes the chemical bonding of a system based on the topology of the quantum charge density. In addition to bonding, AIM allows the calculation of certain physical properties on a per-atom basis, by dividing space up into atomic volumes containing exactly one nucleus. Developed by Professor Richard Bader since the early 1960s, during the past decades QTAIM has gradually become a theory for addressing possible questions regarding chemical systems, in a variety of situations hardly handled before by any other model or theory in Chemistry [1][2]. In QTAIM an atom is defined as a proper open system, i.e. a system that can share energyand electron density, which is localized in the 3D space. Each atom acts as a local attractor of the electron density, and therefore it can be defined in terms of the local curvatures of the electron density. The mathematical study of these features is usually referred in the literature as charge density topology. Nevertheless, the term topology is used in a different sense in Mathematics 38
    39. 39. 39
    40. 40. • • In chemistry and physics, atomic theory is a theory of the nature of matter,  which states that matter is composed of discrete units called atoms, as opposed  to the obsolete notion that matter could be divided into any arbitrarily small  quantity. It began as a philosophical concept in ancient Greece and India and  entered the scientific mainstream in the early 19th century when discoveries in  the field of chemistry showed that matter did indeed behave as if it were made  up of particles. The word "atom" (from the ancient Greek adjective atoms', 'undivisible'[1]) was  applied to the basic particle that constituted a chemical element, because the  chemists of the era believed that these were the fundamental particles of matter.  However, around the turn of the 20th century, through various experiments  with electromagnetism and radioactivity, physicists discovered that the so-called  "indivisible atom" was actually a conglomerate of various subatomic particles ( chiefly, electrons, protons and neutrons) which can exist separately from each  other. In fact, in certain extreme environments such as neutron stars, extreme  temperature and pressure prevents atoms from existing at all. Since atoms were  found to be actually divisible, physicists later invented the term " elementary particles" to describe indivisible particles. The field of science which  studies subatomic particles is particle physics, and it is in this field that  physicists hope to discover the true fundamental nature of matter. 40
    41. 41. Earliest empirical evidence • • Near the end of the 18th century, two laws about chemical reactions  emerged without referring to the notion of an atomic theory. The first was  the law of conservation of mass, formulated by Antoine Lavoisier in 1789,  which states that the total mass in a chemical reaction remains constant  (that is, the reactants have the same mass as the products).[2] The second  was the law of definite proportions. First proven by the French chemist  Joseph Louis Proust in 1799,[3] this law states that if a compound is broken  down into its constituent elements, then the masses of the constituents will  always have the same proportions, regardless of the quantity or source of  the original substance. John Dalton studied and expanded upon this previous work and developed  the law of multiple proportions: if two elements can together form more than  one compound, then the ratios of the masses of the second element which  combine with a fixed mass of the first element will be ratios of small integers . For instance, Proust had studied tin oxides and found that their masses  were either 88.1% tin and 11.9% oxygen or 78.7% tin and 21.3% oxygen  (these were tin(II) oxide and tin dioxide respectively). Dalton noted from  these percentages that 100g of tin will combine either with 13.5g or 27g of  oxygen; 13.5 and 27 form a ratio of 1:2. Dalton found an atomic theory of  matter could elegantly explain this common pattern in chemistry - in the  case of Proust's tin oxides, one tin atom will combine with either one or two  oxygen atoms.[4] 41
    42. 42. • • • • Dalton also believed atomic theory could explain why water  absorbed different gases in different proportions: for example, he  found that water absorbed carbon dioxide far better than it  absorbed nitrogen.[5] Dalton hypothesized this was due to the  differences in mass and complexity of the gases' respective  particles. Indeed, carbon dioxide molecules (CO2) are heavier and  larger than nitrogen molecules (N2). Dalton proposed that each chemical element is composed of atoms  of a single, unique type, and though they cannot be altered or  destroyed by chemical means, they can combine to form more  complex structures (chemical compounds). This marked the first  truly scientific theory of the atom, since Dalton reached his  conclusions by experimentation and examination of the results in an  empirical fashion. Various atoms and molecules as depicted in John Dalton's A New System of Chemical Philosophy (1808). In 1803 Dalton orally presented his first list of relative atomic  weights for a number of substances. This paper was published in  1805, but he did not discuss there exactly how he obtained these  figures.[5] The method was first revealed in 1807 by his acquaintance  Thomas Thomson, in the third edition of Thomson's textbook, A System of Chemistry. Finally, Dalton published a full account in his  own textbook, A New System of Chemical Philosophy, 1808 and  1810. 42
    43. 43. • • • Dalton estimated the atomic weights according to the mass ratios in which they  combined, with hydrogen being the basic unit. However, Dalton did not conceive  that with some elements atoms exist in molecules — e.g. pure oxygen exists as  O2. He also mistakenly believed that the simplest compound between any two  elements is always one atom of each (so he thought water was HO, not H 2O).[6]  This, in addition to the crudity of his equipment, resulted in his table being  highly flawed. For instance, he believed oxygen atoms were 5.5 times heavier  than hydrogen atoms, because in water he measured 5.5 grams of oxygen for  every 1 gram of hydrogen and believed the formula for water was HO (an  oxygen atom is actually 16 times heavier than a hydrogen atom). The flaw in Dalton's theory was corrected in 1811 by Amedeo Avogadro.  Avogadro had proposed that equal volumes of any two gases, at equal  temperature and pressure, contain equal numbers of molecules (in other words,  the mass of a gas's particles does not affect its volume).[7] Avogadro's law  allowed him to deduce the diatomic nature of numerous gases by studying the  volumes at which they reacted. For instance: since two liters of hydrogen will  react with just one liter of oxygen to produce two liters of water vapor (at  constant pressure and temperature), it meant a single oxygen molecule splits in  two in order to form two particles of water. Thus, Avogadro was able to offer  more accurate estimates of the atomic mass of oxygen and various other  elements, and firmly established the distinction between molecules and atoms. In 1827, the British botanist Robert Brown observed that pollen particles floating  in water constantly jiggled about for no apparent reason. In 1905,  Albert Einstein theorized that this Brownian motion was caused by the water  molecules continuously knocking the grains about, and developed a hypothetical  mathematical model to describe it.[8] This model was validated experimentally in  1908 by French physicist Jean Perrin, thus providing additional validation for  particle theory (and by extension atomic theory). 43
    44. 44. Discovery of subatomic particles • • Atoms were thought to be the smallest possible division of matter until 1897  when J.J. Thomson discovered the electron through his work on  cathode rays.[9] A Crookes tube is a sealed glass container in which two  electrodes are separated by a vacuum. When a voltage is applied across  the electrodes, cathode rays are generated, creating a glowing patch where  they strike the glass at the opposite end of the tube. Through  experimentation, Thomson discovered that the rays could be deflected by  an electric field (in addition to magnetic fields, which was already known).  He concluded that these rays, rather than being a form of light, were  composed of very light negatively charged particles he called "corpuscles"  (they would later be renamed electrons by other scientists). Thomson believed that the corpuscles emerged from the molecules of gas  around the cathode. He thus concluded that atoms were divisible, and that  the corpuscles were their building blocks. To explain the overall neutral  charge of the atom, he proposed that the corpuscles were distributed in a  uniform sea of positive charge; this was the plum pudding model[10] as the  electrons were embedded in the positive charge like plums in a plum  pudding (although in Thomson's model they were not stationary) 44
    45. 45. First steps toward a quantum physical model of the atom • • • • The planetary model of the atom had two significant shortcomings. The first is that,  unlike planets orbiting a sun, electrons are charged particles. An accelerating  electric charge is known to emit electromagnetic waves according to the Larmor formula in classical electromagnetism; an orbiting charge should steadily lose  energy and spiral toward the nucleus, colliding with it in a small fraction of a second.  The second problem was that the planetary model could not explain the highly  peaked emission and absorption spectra of atoms that were observed. The Bohr model of the atom Quantum theory revolutionized physics at the beginning of the 20th century, when  Max Planck and Albert Einstein postulated that light energy is emitted or absorbed in  discrete amounts known as quanta (singular, quantum). In 1913, Niels Bohr  incorporated this idea into his Bohr model of the atom, in which an electron could  only orbit the nucleus in particular circular orbits with fixed angular momentum and  energy, its distance from the nucleus (i.e., their radii) being proportional to its energy. [13]  Under this model an electron could not spiral into the nucleus because it could not  lose energy in a continuous manner; instead, it could only make instantaneous " quantum leaps" between the fixedenergy levels.[13] When this occurred, light was  emitted or absorbed at a frequency proportional to the change in energy (hence the  absorption and emission of light in discrete spectra).[13] Bohr's model was not perfect. It could only predict the spectral lines of hydrogen; it  couldn't predict those of multielectron atoms. Worse still, as spectrographic  technology improved, additional spectral lines in hydrogen were observed which  Bohr's model couldn't explain. In 1916, Arnold Sommerfeld added elliptical orbits to  the Bohr model to explain the extra emission lines, but this made the model very  difficult to use, and it still couldn't explain more complex atoms. 45
    46. 46. Discovery of isotopes • While experimenting with the products of radioactive decay, in  1913 radio chemist Frederick Soddy discovered that there  appeared to be more than one element at each position on  the periodic table.[14] The term isotope was coined by Margaret  Todd as a suitable name for these elements. • That same year, J.J. Thomson conducted an experiment in  which he channeled a stream of neon ions through magnetic  and electric fields, striking a photographic plate at the other  end. He observed two glowing patches on the plate, which  suggested two different deflection trajectories. Thomson  concluded this was because some of the neon ions had a  different mass.[15] The nature of this differing mass would later  be explained by the discovery of neutrons in 1940. 46
    47. 47. Quantum physical models of the atom • n 1924, Louis de Broglie proposed that all moving particles — particularly  subatomic particles such as electrons — exhibit a degree of wave-like  behavior. Erwin Schrödinger, fascinated by this idea, explored whether or  not the movement of an electron in an atom could be better explained as a  wave rather than as a particle. Schrödinger's equation, published in 1926, [18]  describes an electron as a wavefunction instead of as a point particle.  This approach elegantly predicted many of the spectral phenomena that  Bohr's model failed to explain. Although this concept was mathematically  convenient, it was difficult to visualize, and faced opposition.[19] One of its  critics, Max Born, proposed instead that Schrödinger's wave function  described not the electron but rather all its possible states, and thus could  be used to calculate the probability of finding an electron at any given  location around the nucleus.[20]This reconciled the two opposing theories of  particle versus wave electrons and the idea of wave-particle duality was  introduced. This theory stated that the electron may exhibit the properties of  both a wave and a particle. For example, it can be refracted like a wave, and  has mass like a particle.[21] 47
    48. 48. • A consequence of describing electrons as  waveforms is that it is mathematically impossible to  simultaneously derive the position and momentum  of an electron; this became known as the  Heisenberg uncertainty principle after the theoretical  physicist Walter Heisenberg, who first described it.  This invalidated Bohr's model, with its neat, clearly  defined circular orbits. The modern model of the  atom describes the positions of electrons in an atom  in terms of probabilities. An electron can potentially  be found at any distance from the nucleus, but,  depending on its energy level, tends to exist more  frequently in certain regions around the nucleus  than others; this pattern is referred to as its atomic  orbital. The orbital's come in a variety of shapes,  manifesting from a simple sphere of the full helium  orbital, to the dumbbell shape of the full neon orbital,  with the nucleus in the middle.[22] 48
    49. 49.                                                                                                                                              Atoms 49
    50. 50. THANKS To:- Mrs. Samnol  PRESENTED BY :SHIVAM CLASS :- IX-B 50