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CHAPTER 2 ATOMS, MOLECULES AND IONS Over the centuries people have often wondered what matter is. Now, most agree that it is composed of: Atoms The smallest particle that can be obtained chemically and is recognizable as an element (and is in turn composed of neutrons, protons and electrons). Molecules Many substances are composed of identical aggregates of atoms we call molecules, eg. H2O, CO2. Ions Some substances contain charged species called ions. The ions may be charged atoms or molecules but overall the charges cancel to give a neutral substance. NaCl, Na+Cl–, NaNO3 , Na+NO3– Some “oft met” ions: Cations (+ve) Anions(-ve) Na+ sodium ion Cl– chloride ion Ca2+ calcium ion NO3– nitrate ion Al3+ aluminium ion SO42– sulfate ion CH3COO– ethanoate (acetate) ion NH4+ ammonium ion 2:1 Daltons Atomic Theory John Dalton (in 1808) An element is composed of tiny particles called atoms. Atoms of the same element have identical properties that are different to those of the atoms of other elements. In a chemical reaction atoms are neither created, destroyed or charged into other atoms. Compounds are formed when two or more elements combine. When these postulates are combined they explain the earlier postulates that had been developed: The Law of Constant Composition. The Law of Conservation of Mass Mass is neither created nor destroyed in chemical reactions. Proust’s Law(1799) Law of Definite Proportions. Different samples of the same chemical substance always contain the same proportion of elements by mass. The Law of Multiple Proportions If two elements can combine in different ways to form different substances, the mass ratio are small ,whole number multiples of each other. The Law of Multiple Proportions Where elements combine to form more than one compound they combine so that a fixed mass of one combines with masses of the other in simple whole number proportions. CO C 12 g CO2 O 16 g C 12 g 1:1 O 32 g 1:2 12 g C 16 g O Mass ratio in CO = 12 g C = 2 Mass ratio in CO 2 32 g O 2:2 Structure of the Atoms The Electron When electricity is passed through a gas at low pressure, cathode rays are produced. Thomson (1856 - 1940) in 1900 showed that the cathode rays consisted of particles that were deflected by both magnetic and electric fields and of much lower mass than that of least heavy atom (hydrogen, H) - so they are subatomic particles. He called them electrons. Later the mass was calculated (R. 1 A. Millikan) and shown to be about 1840 of the mass of the H atom or 9.109390 x 10-28 g. charge on electron e– mass of electron -1 about 1840 of H atom 1 “Prodding atoms with sticks or firing cannon balls at paper? 2:3 Rutherford (1910) bombarded a gold foil with α-particles (He2+ nuclei from the radioactive element radium, Ra). The α-particles were known to be about 7000 times as heavy as the electron and their path was not expected to be affected much by a thin gold foil. The surprise was that while most of the rays were hardly affected by the gold foil (the expected result) a few were strongly deflected and sometimes even reflected. The only way that the observations could be explained was that α-particles had come near some small, but massive, positively charged particle that he called the nucleus. Calculations showed that the nucleus must contain most of the mass of the atom and must be very small compared to the volume occupied by the atom. The positively charged particle present in the nucleus was called a proton. The nucleus of the hydrogen atom carries one positive charge and is a proton. The nuclei of other atoms contain more than one proton. Unfortunately it was nor possible to account for all of the mass of the nuclei of other atoms there appearing to be some additional particle present. Rutherford called these neutrons. Neutrons were observed a few years later. protons mass/u 1.00728 charge +1 neutrons 1.00867 0 nucleus Atomic Number The atomic number is the number of protons in the nucleus and defines the element. All atoms with a particular atomic number are atoms of the same element. Z atomic number 23 Na 11 mass number A For a neutral atom the number of protons must equal number of electrons. The nucleus of a hydrogen atom is a proton. Using sodium as an example: charge Na 12 neutrons 0 11 protons +11 11 electrons -11 0 2:4 total All atoms of an element have the same atomic number, Z. However, it soon became clear that not all have the same mass number, A. Atoms with the same atomic number, Z, but different mass number, A, are called ISOTOPES. The two natural isotopes of chlorine. chlorine -35 35 17 Cl 17 p + 18 n chlorine -37 37 17 Cl 17 p + 20 n The Periodic Table In the early nineteenth century chemists became interested in the physical and chemical similarities that exist between elements. Dobereiner (1817 and 1829) published his ideas about the existence of triads of similar elements. Newlands (1863 - 1866) developed his “law of octaves” . Newlands arranged the element in order of increasing atomic weight and recognized that the fist element was similar to the eighth, the second to the ninth and so on. Mendeleev periodic table. Todays periodic table lists the element in order of increasing atomic number (not atomic weight/mass) GROUPS or FAMILIES PERIODS H Li Be B C N O F The periodic table in the front of your text book has additional columns inserted. The GROUPS have been numbered in various ways but we will use the modern system of numbering 1 - 18 or old 1 - 8. (Make sure you understand the system) We sub-divide the periodic table into: Main Groups - Groups 1, 2 and 13 to 18 Transition Metals 4 to 11 or d-block (3-12) (explained later) 2:5 He Ne Metals on the left and non-metals on the right Metals Metalloids Nonmetals Some of the groups have names in common usage Group 1 Group 2 Group 16 Group 17 Group 18 alkali metals alkaline earths Oxygen family halogens noble gases Molecules and Ions Molecules Molecular formula Smallest complete unit of molecular substance carbon dioxide is written as CO2. The number after the symbol tells how many of that kind of atom is present in the molecule.. sulfur trioxide SO3; methanol CH4O; nitric acid HNO3 (in aqueous solution). These are the formulas and tell little about the real structure of the molecule. H The structural formula for methanol is H C O H and this tells how the atoms are connected. H Nearly the same information can be conveyed with the condensed molecular formula CH3OH - 3 H connected to the C. O connected to the C and an H connected to the O. A ball and stick model gives a much better representation. 2:6 Structural formula H H C O H O O C O S O O propanone "acetone" H H C H C O H C H H Sometimes the atoms can be connected together in several different ways. The compound represented are different and called ISOMERS. ethanol C2H6O H H H H H C O C H H H H C C O H H H Structural isomers CH3CH2OH CH3OCH3 condensed formulas Monatomic ions 23 + 11 Na 23 11 Na + e− 11p+ + 10e– 11p+ + 11e– cation 16 8O + 2 16 2− 8O e– 8p+ + 8e– 8p+ + 10e– anion Monatomic Ions 13 14 15 Group I Group II Group III Group IV Group V +1 +2 Al3+ - -3 2:7 16 17 Group VI Group VII -2 -1 18 Group 0 0 If you count the electrons you will find that when the ions form, they usually have the same number of electron as the previous or next noble gas. This is not always the case and some ions (monatomic) may exhibit more than one charge. When this is the case it is necessary to distinguish between the ions using the roman numeral after the name to indicate the particular ion intended. chromium(III), Cr3+ iron(III), Fe3+ copper(II), Cu2+ chromic ferric cupric chromium(II), Cr2+ iron(II), Fe2+ copper(I), Cu+ chromous ferrous cuprous chromium iron copper Anions (See page 45) chlorine gives chloride ion, Cl– etc. Polyatomic ions For example NO −3 nitrate ion NO −2 nitrous ion OH– hydroxide ion NH +4 ammonium ion Learn these - See Table 2.2 and 2.3 Th d oxoanions. 2:8 Molecular Elements Some elements in the normal state at room temperature are molecular. Examples are nitrogen, N2; oxygen, O2; the halogens, eg. Cl2. Nomenclature Binary Compounds Binary compounds are easy to name. The name consists of two parts: NaCl name of element part of name of element + IDE sodium chloride HF hydrogen fluoride CaO calcium oxide Al2O3 aluminum oxide All have the name ending IDE. There is a problem where several different binary compounds are formed by the pair of elements. In this case the element names are prefixed by a Greek multiplier prefix. mono tetra hepta deca 1 4 7 10 di penta octa undeca 2 5 8 11 tri hexa nona dodeca 3 6 9 12 So, SCl2 is sulfur dichloride, SCl4 is sulfur tetrachloride and S2Cl2 is disulfur dichloride. Some binary compounds have common names that replace the systematic names. water, H2O; ammonia, NH3; methane, CH4. There are more examples in the textbook. Acids You will meet a group of compounds called acid as you do your labs. They are usually only correctly described as acids in aqueous solutions in which they are more or less ionized. 2:9 Pure Substance HCl(g) hydrogen chloride HBr(g) hydrogen bromide HI(g) hydrogen iodide HNO3 HNO2 Aqueous solution H+(aq), Cl–(aq) hydrochloric acid H+(aq), Br–(aq) hydrobromic acid H+(aq), I–(aq) hydriodic acid Oxoanions hydrogen nitrate H+(aq), NO3–(aq) nitric acid H+(aq), NO2–(aq) nitrous acid hydrogen nitrite etc. see page 56 of your text Naming Ionic Compounds Ionic compounds are named cation name first then anion name. Thus KNO3(s) Ca(NO3)2(s) FeSO4(s) Fe2(SO4)3(s) potassium nitrate calcium nitrate iron(II) sulfate iron(III) sulfate Practice naming compounds. 2:10 Anion name chloride bromide iodide nitrate nitrite