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Unit 2 Atoms, Molecules, & Ions Law of Constant Composition (Definite Proportion): All samples of a compound have the same composition, the same proportion by mass of the constituent elements. H2O by mass is always 11.19 % H 88.81 % O A 0.1 g sample of Mg, when combined with O2 (g) yields 0.166 g MgO. A second sample, 0.144 g of Mg is also combined with O2 (g). What is the mass of MgO produced? Law of Multiple Proportions: When 2 elements form a series of compounds, the ratio of the masses of the second element that combine with 1.0 g of the first element can always be reduced to small whole numbers. The following data were collected for several compounds of nitrogen combining with 1 g oxygen. Determine the formulas. Compound A = 1.750 g N Compound B = 0.8750 g N Compound C = 0.4375 g N Dalton’s Atomic Theory Each element is made up of tiny particles called atoms. The atoms of a given element are identical; atoms of different elements are different in some fundamental way. Chemical compounds are formed when atoms of different elements combine with each other. They combine in simple whole number ratio, always with the same relative number and types of atoms. Chemical reactions involve reorganization of the atoms, changes in the way they are bonded together. The atoms themselves are indestructible. History 1803 John Dalton discovered that elements are made of atoms. He thought that atoms were solid, like a marble. Early 1800’s Michael Faraday discovered that a cathode ray is negatively charged particles 1874 George Stoney coined the term “electron” 1875 Crooks discovered the electron has a mass of 9.1 x 10-28 g and have a negative (-) charge. 1897 JJ Thomson used a cathode ray tube to discover negative ray consisting of electrons with a charge to mass ratio c/m = -1.76 x 108 c/g 1895 Roentgen discovered when using a cathode ray tube, that there was fluorescence in the room. He found X rays. 1909 Robert Millikan worked with an oil drop to determine the magnitude of e- charge to be – 1.6022 x 19-19 coulomb. 1896 Henri Bacquerel (1852 – 1908) discovered radioactivity a alpha particle is a helium nucleus b beta particle is an electron g gamma ray is pure energy 4He 2 e- 1911 Earnest Rutherford (1871 – 1937) discovered that there is a mass in the center of the atom with a positive charge containing p+, the e- are around the nucleus, with the atom mostly empty space. 1932 Chadwick discovered the neutron which is also found in the nucleus. Subatomic particles Particle amu mass (g) Atomic charge Electrical charge (coulomb) Proton (p+) 1.0073 1.673 x 10-24 +1 +1.602 x 10-19 Electron (e-) .0005486 9.109 x 10-28 -1 -1.602 x 10-19 Neutron (n) 1.0087 1.675 x 10-24 0 0 Nucleus has a density of 1 x 1013 to 1 x 1014 g/cm3 There are 4 forces that hold an atom together: 1. gravitational – attraction between 2 bodies with mass 2. electromagnetic – attractive and repulsive forces due to charged or magnetic objects 3. strong nuclear force – keep protons which all have positive charges from flying apart 4. weak nuclear force – responsible for (b) beta decay in radioactivity Periodic Table Atomic mass 12.0111 C Element symbol 6 Atomic number Elements with the atomic mass in ( ) means that very little of it exists at any given time. Scientists have not been able to attain an good average atomic mass. Nuclear Notation Atomic Mass Atomic number A Z X Chemical symbol Atomic mass = p+ + N Atomic number = p+ Isotopes = Atoms of the same element (so the same atomic number) but with a different mass because they have a different number of neutrons. Kinds of Decay Alpha decay a A Z X A-4 N+ Z-2 4 He N = new element Beta decay b X Z A N + e Z+1 A N p+ + e- Atomic Mass is the average of the number of protons and Neutrons found in the nucleus Average fraction of abundance ) ( Mass of = S ( isotope Atomic Mass of each isotope each ) In naturally occurring Argon, 99.600% of the atoms are 40Ar with a mass of 39.9624 amu (u), 0.337% 36Ar with 18 18 38 a mass of 35.96755 u, and 0.063% 18Ar with a mass of 37.96272 u. Calculate the Average mass of Argon. Classification of elements as metals, nonmetals, and metalloids. Alkali Earth metals +2 ion Alkali metals +1 ion Rare Earth Metals Noble gases Halogens -1 ion Chalcogens -2 ion Transition metals Periodic table with atomic symbols, atomic numbers, and partial electron configurations. Diatomic molecules There are 7 elements that do not exist as 1 atom alone. As an element, they are always as 2 atoms bonded together as a molecule. Location of the diatomic elements Molecular Compound – molecule with more than 1 type of atom Molecular Formula – shows the actual # of atoms Empirical Formula – shows the relative # of atoms in the smallest possible whole # ratio Structural Formula – shows how the atoms are assembled into the molecule Chemical Bonds – the forces holding atoms together Ionic bonds – transfer of e- resulting in ions attracted to each other Ions – atom or group of atoms which have gained or lost e- (s), therefore carrying an electrical charge. cation (+) positive charged ion anion (-) negative charged ion Covalent bonds – sharing of e- Polar covalent bond – an uneven sharing of e- resulting in a molecule which appears to have opposites charges on different parts of the molecule. Nonpolar covalent bond – an even sharing of e- Naming compounds Inorganic binary compounds type I • one atom (+) cations maintain atom name • one atom (-) anion changes name (-ide) • (+) cation is followed by the (-) anion Inorganic type II – some anions form multiple charges • Use roman numerals to indicate charge example: CuCl – copper (I) chloride, CuCl2 – copper (II) chloride • Or/ higher charge name ends with (-ic) while lower charge name ends with (-ous) example: FeCl3 ferric chloride FeCl2 ferrous chloride Element Latin naming Copper Cu+2 Cuprous, Cu+3 Cupric Iron Fe+2 Ferrous, Fe+3 Ferric Mercury Hg2+2 Mercurous, Hg+2 Mercuric Lead Pb+2 Plumbous, Pb+4 Plumbic Tin Sn+2 Stannous, Sn+4 Stannic Chromium Cr+2 Chromous, Cr+3 Chromic Manganese Mn+2 Manganous, Mn+3 Manganic Cobalt Co+2 Cobaltous, Co+3 Cobaltic Polyatomic – must memorize the names! Oxy anions: elements with oxygen acting as polyatomic • Less than small number of oxygen (hypo-) • Small number of oxygen (-ite) • Large number of oxygen (-ate) • More than large number of oxygen (per-) Example: hypochlorite ClO-1 chlorite ClO2-1 chlorate ClO3-1 perchlorate ClO4-1 Naming Acids – when dissolved in H2O they produce free protons (H+) If anion does not contain oxygen • Use prefix (hydro-) with suffix (-ic) example HS hydrosulfuric acid If anion does contain oxygen • • • If anion ends in (-ate) use root name with (-ic) H2SO4 sulfuric acid If anion ends in (-ite) use root name with (-ous) H2SO3 sulfurous acid When multiple oxyanion use Example HClO4 HClO3 HClO2 HClO perchlorate chlorate chlorite hypochlorite perchloric acid chloric acid chlorous acid hypochlorous acid Binary Covalent type III – 2 nonmetals 1st element in formula named first with full element name 2nd element named as if anion Prefixes used to denote number of atoms present Mono is never used for naming 1st element mono - 1 hexa - 6 di - 2 hepta - 7 tri - 3 octa - 8 tetra - 4 nona - 9 penta - 5 dec - 10 Naming Organic Molecules Name root name according to the number of carbons in the longest continuous chain. Name the suffix according to the bonding, carbon to carbon. • Single bond C – C • Double bond C = C • Triple bond C C end with -ane end with -ene end with -yne Name substitutions on the main chain before the root name, giving the lowest possible number. Hydrocarbon Substitute Naming Group Names Methyl -CH3 Ethyl -CH2CH3 Propyl -CH2CH2CH3 Butyl -CH2CH2CH2CH3 Halides Chloro -Cl Bromo -Br Iodo -I Fluoro -F Prefixes used when more than one group of the same kind is attached ditritetrapentahexa- two substitutes three four five six Naming organic molecules containing functional groups Drop the e, add Functional Group ending -OH Group Name Name Alcohol - ol Carboxylic Acid Ketone - oic Acid - one Functional Group Group Name Aldehyde Name ending - al Drop –ane, add Double bond Triple bond - ene - yne Isomers: Molecules with the same chemical formula, but different arrangement of atoms. With different arrangement of atoms, the properties of the chemicals are different. Small molecules only have one arrangement possible. CH4 Methane C2H6 Ethane C5H12 has 3 isomers