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Transcript
Elements and Compounds Composition of Matter Elements 1 Element Identification 2 Source of Light Elements Solar Spectra 3 Detecting Elements Molecular Spectrum of CO 4 Chemical and Physical Properties • Elements • An Element is the basic form of matter – it cannot be broken down into smaller parts by chemical methods Each element has a unique identity Compounds • A Compound is a unique substance that is formed from two or more elements and has different properties from all other substances 5 Pure Substances • Pure substances have unchanging, definite compositions and properties. • Table Salt: crystalline white substance that is commonly used to modify food taste NaCl : 1 sodium atom + 1 chlorine atom. Two different elements, unique properties. Properties of Substances • A physical property is one that can be observed without performing a change to the substance: • Color • Density • Hardness • etc. 6 Names of the Elements • Latin names from historical context produce unusual symbols for some elements: first letter is capitalized, second is lower case – Ferrum: Fe (Iron) – Aurum: Au (Gold) – Cobalt: Co 7 Atomic Theory • John Dalton used Greek ideas with experimentation to reveal the atomic nature of the elements • The atom is defined as the smallest particle of an element that retains the properties of the bulk element • Daltons atomic theory allows the prediction of composition Daltons Atomic Theory • Matter is composed of atoms • All atoms of each element are identical and possess the same properties • Chemical compounds are composed of atoms of different elements in whole number ratio • Reactions produce re-arranged atoms 8 Law of Constant Composition • Any compound is made up of elements in the same proportion by mass. • A compound is identified by a specific formula Compounds and Molecules • Two or more elements bonded to each other are molecules or molecular compounds • Elements bonded together are polyatomic i.e. Cl2 • Molecules may be formed by both types: compounds or polyatomic elements 9 Chemical Properties • Chemical properties are related to the identity of the element or the compound • The ability of a substance to undergo a chemical change: “reaction” • Chemical change is a transformation of a substance into a new substance Mixtures of elements and compounds 10 Mixtures • Heterogeneous: a non-uniform mixture containing two or more phases • Sample of mixture produces different amounts of each substance in container Homogeneous Mixtures • One physical state with uniform properties • May be a Solution or an Alloy – Solution: homogeneous mixture with one liquid phase – Alloy: homogeneous mixture of solids with one phase (solid solution) 11 Electrons • 1897 J.J. Thompson proves existence of the electron • Is common to all atoms • Has a charge of -1 • Used “Plum Pudding” analogy : electrons and a positive medium mixed as a solid particle Rutherford Experiment • Used positively charged helium ions to bombard thin gold foil • Showed a dense positively charged center- not a conglomeration of positive and negative charges • Calculated charge to mass ratio =-1.76 x 108 coulombs/gram From an electron beam 12 Electron Properties • R.A. Millikan (Univ. Of Chicago) determined charge magnitude of electrons which led to the determination of the mass • q= -1.6 x 10-19 Coulombs • Mass = 9.11 x 10-28 grams (from charge to mass ratio) 13 The Nuclear Atom Counting The Particles • All atoms of each element are not quite the same • Mass of atoms of an element is variable, a result of a new particle, the neutron • The nuclear mass is the sum of protons and neutrons: isotopes 14 Calculating Isotopic Values • Atomic Number = element identity • = Protons in nucleus • neutrons=mass number – protons • 90Sr has 38 protons: • 90 – 38 = 52 neutrons • Still has 38 electrons! atomic number Every atom with an atomic number of 1 is a hydrogen atom. 1 proton in the nucleus 15 atomic number Every atom with an atomic number of 6 is a carbon atom. 6 protons in the nucleus atomic number Every atom with an atomic number of 92 is a uranium atom. 92 protons in the nucleus 16 Isotopes of the Elements • Atoms of the same element can have different masses. • They always have the same number of protons, but they can have different numbers of neutrons in their nuclei. • The difference in the number of neutrons accounts for the difference in mass. • These are isotopes of the same element 17 Isotopes of the Same Element Have Equal numbers of protons Different numbers of neutrons Isotopic Notation 18 Isotopic Notation 6 protons + 6 neutrons 12 6C 6 protons Isotopic Notation 6 protons + 8 neutrons 14 6C 6 protons 19 Isotopic Notation 8 protons + 8 neutrons 16 8O 8 protons Isotopic Notation 8 protons + 9 neutrons 17 8O 8 protons 20 Isotopic Notation 8 protons + 10 neutrons 18 8O 8 protons Hydrogen has three isotopes 1 proton 1 proton 1 proton 0 neutrons 1 neutron 2 neutrons 21 Examples of Isotopes Element Protons Electrons Symbol Neutrons Hydrogen Hydrogen Hydrogen 1 1 1 0 1 2 Uranium 235U 92 Uranium 238U 92 92 92 143 92 92 146 1 1 1 1 1H 1 2H 1 3H Atomic Mass 22 • The mass of a single atom is too small to measure on a balance. • Using a mass spectrometer, the mass of the hydrogen atom was determined. A Modern Mass Spectrometer Positive ions formed from sample. Electrical field From the intensity and positions at slits A mass of the lines Deflection on the mass of accelerates spectrogram spectrogram, the different ions positive ions. positive is recorded. isotopes and their at relative occurs amounts can be determined. magnetic field. 5.8 23 A typical reading from a mass spectrometer. The two principal isotopes of copper are shown with the abundance (%) given. 5.9 Using a mass spectrometer, the mass of one hydrogen atom was determined to be 1.673 x 10-24 g 24 This number is very small. small small small small small small small small small small small small small small small small small small small small small The mass of a hydrogen atom is very small. Numbers of thisthis sizeproblem are too small for of To overcome a system practical use. relative atomic masses using “atomic mass units” was devised to express the masses of elements using simple numbers. 1.673 x 10-24 g 25 The standard to which the masses of all other atoms are compared to was chosen to be the most abundant isotope of carbon. 12 6 A mass of exactly 12 atomic mass units (amu) was assigned to 12 6 26 1 1 amu is defined as exactly equal to 12 the mass of a carbon-12 atom 1 amu = 1.6606 x 10-24 g 12 6 Average atomic mass 1.00797 amu. 27 Average atomic mass 39.098 amu. Average atomic mass 248.029 amu. 28 Average Relative Atomic Mass • Most elements occur as mixtures of isotopes. • Isotopes of the same element have different masses. • The listed atomic mass of an element is the average relative mass of the isotopes of that element compared to the mass of carbon-12 (exactly 12.0000…amu 29 To calculate the atomic mass multiply the atomic mass of each isotope by its percent abundance and add the Average atomic results. mass (amu) Isotope Isotopic mass (amu) Abundance (%) 63 29 Cu 62.9298 69.09 65 29 Cu 64.9278 30.91 63.55 (62.9998 amu) 0.6909 = 43.48 amu (64.9278 amu) 0.3091 = 20.07 amu 63.55 amu Relationship Between Mass Number and Atomic Number 30 The mass number minus the atomic number equals the number of neutrons in the nucleus. mass number atomic number 109 47 atomic number 47 mass number 109 Ag = = number of neutrons 62 31 Atomic Mass Units • Isotopic mass is determined by comparison with a standard: • AMU = 1/12 the mass of a 12C atom • Boron isotopes: – 10B is 0.83442 times the mass of 12C: – Therefore has 10.013 AMU – 11B has 11.009 AMU Table Values for Atomic Mass • Since elements exist in nature with isotopic masses, the reported value is the weighted average of all of these isotopes based on their percent abundance 32 Calculating atomic masses • Nature contains 19.9% 10B and 80.1% 11B. What is the weighted average? • 10B: 0.199 x 10.013 AMU = 1.99 AMU • 11B: 0.801 x 11.009 AMU = 8.82 AMU • 10.81 AMU Binary Compounds • Metal-Nonmetal • Single Oxidation • Multiple oxidation • Polyatomic ions Al2O3 CaCl2 Fe2O3 ; FeCl2 Mg3(PO4)2 33 Practice: Nomenclature • • • • • • • • • • • • NaCl Sodium Chloride HCl (blue Stuff) Hydrogen Chloride H2O (blue Stuff) Dihydrogen oxide FeS Iron(II) oxide Al2O3 Aluminum oxide SiO2 (blue stuff) Silicon dioxide Binary Compounds • • • • • • • • NaCl Sodium chloride CaBr2 Calcium bromide NH4Cl Ammonium chloride N2O Dinitrogen oxide • • • • • • • • FeS Iron(II) sulfide Mn2O3 Manganese(III) oxide Sn(NO2)4 Tin(IV) nitrite Fe(OH)3 Iron(III) hydroxide 34 Non-metal Binary Compounds • Use Greek Prefixes (table 4-3) • Write first element which is closest to metal • Hydrogen is written second • First element written as english name • Second element(s) end with –ide • Never use mono- for first element; only remaining element Oxides of Nitrogen • • • • • • • • N2O Dinitrogen monoxide NO Nitrogen monoxide N2O3 Dinitrogen trioxide N2O5 Dinitrogen pentoxide 35 Binary Acids • Hydrogen plus halogen (group VII) or other anions • Name of compound ends with –ic acid • HCL – hydrogen choride • Hydrochloric acid • HBr- hydrogen bromide • Hydrobromic acid Oxy Acids • Formed by the combination of hydrogen and most polyatomic ions • Name after the root of the anion plus the following endings: • -ic acid if the polyatomic ion ended with – ate • -ous if the anion ended with –ite • Mineral acids use common names 36 Oxyacids • • • • • • • • C2H3O2- acetate ion CO32- carbonate ion NO3- nitrate ion PO43- phosphate ion ClO- hypochlorite ion ClO2- chlorite ion ClO3- chlorate ion ClO4- perchlorate ion • • • • • • • • HC2H3O2 acetic acid H2CO3 carbonic acid HNO3 nitric acid H3PO4 phosphoric a. HClO hypochlorous a. HClO2 chlorous acid HClO3 chloric acid HClO4 perchloric acid 37