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Atoms 18 2 Chapter Atoms 2.1 The Atomic Theory of Matter Chemists make their observations in the macroscopic world and seek to understand the fundamental properties of matter at the level of the microscopic world (i.e. molecules and atoms). The reason why certain chemicals react the way they do is a direct consequence of their atomic structure. The word "atom" is derived from the Greek word "atomos", meaning indivisible. The philosopher Democritus (460-370 B.C.) believed that matter was composed of fundamentally indivisible particles, called "atomos". Dalton's atomic theory of 1803: 1. Each element is composed of extremely small indivisible particles called atoms. 2. All atoms of a given element are identical; the atoms of different elements are different and have different properties (including different masses). 3. Atoms of an element are not changed into different types of atoms by chemical reactions; atoms are neither created nor destroyed in chemical reactions, they simply combine, separate or rearrange. 4. Compounds are formed when atoms of more than one element combine in a specific simple ratio; a given compound always has the same relative number and kind of atoms. Atoms are the basic building blocks of matter; they are the smallest units of an element: Atoms are the smallest particle of an element which retains the chemical properties of that element Simple "laws" (i.e. theories) of chemical combination which were known at the time of Dalton: 1. The law of constant composition (in a given compound the relative number and kind of atoms are constant). For example, water is made up of hydrogen and oxygen in the fixed ratio of H:O 2:1, or a fixed mass ratio of 11% to 89%. 2. The law of conservation of mass (the total mass of materials present after a chemical reaction is the same as the total mass before the reaction). In a chemical reaction mass of reactants = mass of products (Lavoisier’s law) Atoms 19 2.2 The Discovery of Atomic Structure 1803 Dalton - the atom is a indivisible, indestructible, tiny ball 1850 Evidence is accumulating that the atom is itself composed of smaller particles The current model... Cathode rays and electrons Electrical discharge through partially evacuated tubes produced radiation. This radiation originated from the negative electrode, known as the cathode (thus, these rays were termed cathode rays). The "rays" traveled towards, or were attracted to the positive electrode (anode) Not directly visible but could be detected by their ability to cause other materials to glow, or fluoresce Traveled in a straight line Their path could be "bent" by the influence of magnetic or electrical fields A metal plate in the path of the "cathode rays" acquired a negative charge The "cathode rays" produced by cathodes of different materials appeared to have the same properties These observations indicated that the cathode ray radiation was composed of negatively charged particles (now known as electrons). J.J. Thompson (1897) used results from cathode ray tube (commonly abbreviated CRT) experiments to discover the electron. He was able to measure the charge to mass ratio for a stream of electrons (using a cathode ray tube apparatus) at 1.76 x 108 coulombs/gram. The side image is of J.J. Thomson and a cathode ray tube from around 1897, the year he announced the discovery of the electron. Only the end of the CRT can be seen to the right-hand side of the picture. The image below is of a CRT used by Thomson in his experiments. It is about one meter in length and was made entirely by hand. Atoms 20 The amount the cathode ray bent from the straight line using either the electric field or the magnetic field allowed Thomson to calculate the e/m ratio, but he was not able to determine the mass of the electron. However, from his data, if the charge of a single electron could be determined, then the mass of a single electron could be determined. Robert Millikan (1909) was able to successfully measure the charge on a single electron (the "Millikan oil drop experiment"). This value was determined to be 1.60 x 10-19 coulombs. Thus, the mass of a single electron was determined to be: (1 gram/1.76 x 108 coulombs)*(1.60 x 10-19 coulombs) = 9.10 x 10-28 grams Note: the currently accepted value for the mass of the electron is 9.10939 x 10-28 grams. The nuclear atom J.J. Thompson model of the atom (1900) The atom consists of a sphere of positive charge within which was buried negatively charged electrons Also known as the "plum pudding" model of the atom Rutherford model of the atom (1911) (the gold-foil experiment...image aside) Observations Most of the alpha particles pass straight through the gold foil. Some of the alpha particles get deflected by very small amounts. A very few get deflected greatly. Even fewer get bounced of the foil and back to the left. Atoms 21 Conclusions The atom is 99.99% empty space. The nucleus contains a positive charge and most of the mass of the atom. The nucleus is approximately 100,000 times smaller than the atom. Rutherford (1919) discovers protons - positively charged particles in the nucleus Chadwick (1932) discovers neutron - neutral charge particles in the nucleus. We will discuss Bohr’s model and the quantum model later 2.3 Modern view of atomic structure Physicists have identified a long list of particles which make up the atomic nucleus. Chemists, however, are primarily concerned with the following sub-atomic particles: electron proton neutron Electron The electron is negatively charged, with a charge of -1.602 x 10-19 Coulombs (C). For convenience, the charge of atomic and sub-atomic particles are usually described as a multiple of this value (also known as the electronic charge). Thus, the charge of the electron is usually simply referred to as -1. Proton The proton has a charge of +1 electron charge (or, +1.602 x 10-19 C) Atoms 22 Neutron Neutrons have no charge, they are electrically neutral. Note: Because atoms have an equal number of electrons and protons, they have no net electrical charge. Protons and neutrons are located in the nucleus (center) of the atom. The nucleus is small compared to the overall size of the atom. The majority of the space of an atom is the space in which the electrons move around. Electrons are attracted to the protons in the nucleus by the force of attraction between particles of opposite charge. Note: The strength of attraction between electrons and protons in the nuclei for different atoms is the basis of many of the unique properties of different atoms. The electrons play a major role in chemical reactions. In atomic models, the electrons are represented as a diffuse electron cloud The mass of an atom is extremely small. The units of mass used to describe atomic particles is the atomic mass unit (or amu). An atomic mass unit is equal to 1.66054 x 10-24 grams. How do the different sub-atomic particles compare as far as their mass? Proton = 1.0073 amu Neutron = 1.0087 amu Electron = 5.486 x 10-4 amu From this comparison we can see that: The mass of the proton and neutron are nearly identical The nucleus (protons plus neutrons) contains virtually all of the mass of the atom The electrons, while equal and opposite in charge to the protons, have only 0.05% the mass The size of an atom is quite small also, the typical range for atomic diameters is between 1 x 10-10 and 5 x 10-10 meters. Atoms 23 Note: A convenient unit of measurement for atomic distances is the angstrom (Å). The angstrom is equal to 1 x 10-10 meters. Thus, most atoms are between 1 and 5 angstroms in diameter. Parts of an Atom Symbol Mass in amu Relative Electric Charge Proton p 1.0073 +1 Nucleus Neutron n 1.0087 0 Nucleus Electron e 0.0005 −1 Orbital/Electron Cloud Subatomic Particle Atomic Location Isotopes, Atomic Numbers and Mass Numbers What characteristic feature of sub-atomic particles distinguishes one element from another? All atoms of an element have the same number of protons in the nucleus Since the net charge on an atom is 0, the atom must have an equal number of electrons. What about the neutrons? Although usually equal to the number of protons, the number of neutrons can vary somewhat. Atoms which differ only in the number of neutrons are called isotopes. Since the neutron is about 1.0087 amu (the proton is 1.0073), different isotopes have different masses. Carbon All atoms of the element Carbon (C) have 6 protons and 6 electrons. The number of protons in the carbon atom are denoted by a subscript on the left of the atomic symbol: This is called the atomic number, Z, and since it is always 6 for carbon, it is somewhat redundant and usually omitted. Another number, the "Mass Number", A, is a superscript on the left of the atomic symbol. It denotes the sum of the number of protons and neutrons in the particular isotope being described. For example: Refers to an isotope of carbon which has (as expected for the element carbon) six protons, and six neutrons. The following isotope of carbon: Atoms 24 Has 6 protons (atomic number) and 8 neutrons (8=14-6 or A-Z). This isotope is also known simply as "carbon 14". Carbon 12 is the most common form of carbon (~99% of all carbon). An atom of a specific isotope is called a nuclide. Hydrogen, for example has three isotopes: number of protons number of neutrons Mass in amu Hydrogen (protium) Heavy hydrogen (deuterium) 1 1 0 1 1 2 Super-heavy hydrogen (tritium) 1 2 3 Isotope Tritium, however, is unstable and the nucleus decays (breaks apart). This is an example of radioactivity. Since all atoms are composed of protons, electrons and neutrons, all chemical and physical differences between elements are due to the differences in the number of these sub-atomic particles. Therefore, an atom is the smallest sample of an element, because dividing an atom further (into sub-atomic particles) destroys the element's unique identity. Exercises Complete the table by filling in the empty boxes. Element Calcium (Ca) Nickel (Ni) Gold (Au) Atomic Number 79 6 Protons Neutrons 20 28 20 Mass number 59 118 14 Write the symbol or name for each of the following elements, as appropriate. a. b. c. d. e. f. g. h. i. gold Hg uranium sodium Mn Zn fluorine Fe lead Isotope Symbol Atoms 25 2.4 The Periodic Table As more and more elements were discovered and characterized, efforts were made to see whether they could be grouped, or classified, according to their chemical behavior. This effort resulted, in 1869, in the development of the Periodic Table. Certain elements show similar characteristics: Lithium (Li), Sodium (Na) and Potassium (K) are all soft, very reactive metals Helium (He), Neon (Ne) and Argon (Ar) are very non-reactive gasses If the elements are arranged in order of increasing atomic number, their chemical and physical properties are found to show a repeating, or periodic pattern.. As an example of the periodic nature of the atoms (when arranged by atomic number), each of the soft reactive metals comes immediately after one of the nonreactive gasses. The elements in a column of the periodic table are known as a family or group. The labeling of the families are somewhat arbitrary, but are usually divided into the general groups of: Metals (everything on the left and middle region) Non-metals (upper diagonal on the right hand side - green, salmon and red) Metalloids (atoms in the boundary between the metals and metalloids: Boron(B), Silicon(Si), Germanium(Ge), Arsenic(As), Antimony(Sb), Tellurium(Te), and maybe Astatine(At)). These are some of the more useful materials for semi-conductors. Another convention is the 'A' and 'B' designators with column number labels (either in Roman or Arabic numerals). These columns have different types of classifications: Group 1A or I 2A or II 6A or VI 7A or VII 8A or VIII Name Alkali metals Alkaline earth metals Chalcogens ("chalk formers") Halogens ("salt formers") Noble gases (or inert, or rare gases) Elements Li, Na, K, Rb, Cs, Fr Be, Mg, Ca, Sr, Ba, Ra O, S, Se, Te, Po F, Cl, Br, I, At He, Ne, Ar, Kr, Xe, Rn 26 Atoms Group I 1 2 3 4 5 6 7 II III IV V VI VII VIII 1 2 H He 3 4 5 6 7 8 9 10 Li Be B C N O F Ne 11 12 13 14 15 16 17 18 Na Mg Al Si P S Cl Ar 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 87 88 89 104 105 106 107 108 109 110 Fr Ra Ac Rf Db Sg Bh Hs Mt Ds 58 59 60 61 62 63 64 65 66 67 68 69 70 71 Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 90 91 92 93 94 95 96 97 98 99 100 101 102 103 Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Legend Solid Liquid Gas Synthetic Alkali metals Alkali earth metals Transition metals Rare earth metals Other metals Noble gases Halogens Other nonmetals The elements in a family of the periodic table have similar properties because they have the same type of arrangement of electrons at the periphery of their atoms. The majority of elements are metals. Some properties of metals are: high luster high electrical conductivity high heat conductivity solid at room temperature (except Mercury [Hg]) ductile malleable Note: hydrogen is a non-metal (at left hand side of the periodic table) Non-metals can be solid, liquid or gas at room temperature. Their properties are opposite those of metals Atoms EXERCISES Q.1 Give the correct terms that correspond to the following statements: a. Six elements falling on the stair-step line in the Periodic Table. b. The total number of periods in the current Periodic Table. c. Name given to elements of group II. d. Name given to elements that are assigned the columns 3 up to 12. e. Very reactive elements that lose one electron easily. f. Nonmetals with eight valence electrons. g. The person accredited with the discovery of the atomic number. Q.2 Answer with True/False and correct the false statements properly: a. Most metals are solids at room temperature. b. Nonmetals have the general property of being lustrous. c. All the halogens are nonmetals. 27