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Chapter 2: Atoms, Molecules, Ions 2.1 Atomic Theory • Democritus (5th century BC) - matter is composed of indivisible particles call atomos • Alchemists (100 - 1600 A. D.) - made many discoveries while pursuing metaphysical goals • Robert Boyle (17th century) - quantitative measurements of the properties of matter • Joseph Priestley and Antoine-Laurent Lavoisier (18th century) - showed that combustion is a reaction between matter and oxygen • John Dalton - Atomic Theory (ca. 1808) 1. Elements are composed of extremely small particles call atoms. All atoms of a given element are identical, having the same size, mass, and chemical properties. The atoms of one element are different from the atoms of all other elements. (Fig.2.1, p. 36) 2. Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either and integer or a simple fraction. (incorporates laws of definite proportions and multiple proportions) 3. • Law of definite proportions - different samples of the same compound always contain its constituent elements in the same proportions by mass. • Law of multiple proportions - If two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers (use CO and CO2 as an example: 12 to 16 versus 12 to 32.) • Law of conservation of mass - matter can be neither created nor destroyed. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction. 2 2.2 The Structure of an Atom An atom is the basic unit of an element that can enter into chemical combinations. But what is an atom made of: Order of discovery: Cathode Ray tube (Figure 2.3) charge to mass ratio -1.76 × 108 coulomb/g Millikan Oil drop experiment • charge of electron = -1.60 × 10-19 coulombs • mass of the electron is 9.09 × 10-28 g (very small !) • Atoms are composed of • electrons negatively charged mass = 9.1096 × 10-31 kg • protons positively charged mass = 1.6727 × 10-27 kg • neutrons neutral mass = 1.6750 × 10-27 kg • Radioactivity - Radioactive substances emit 3 kinds of particles and/or radiation • alpha (α) particles (helium nuclei) • beta (β) particles (electrons) • gamma (γ) rays Radioactive Particles (Fig. 2.6) Rutherford experiment (Fig. 2.8) relate to α particles • Proton Chapter 2 Chem 1303 3 • Proton Ernest Rutherford discovered that the positive charge was not spread out over the entire volume of the atom but was concentrated in a small, very dense core (the nucleus). Have the same but opposite charge as the electrons, but 2000 times more mass • Neutron discovered by James Chadwick in 1932 2.3. Atomic Number, Mass Number, and Isotopes • Atomic number (Z) = number of protons in nucleus = number of electrons in a neutral atom. • Mass number (A) = sum of number of protons plus number of neutrons in nucleus A Z X Examples of symbols: H, C, O, U-235 (92) • Element - a form of matter in which all of the atoms have the same atomic number • Two atoms of the same element can have different mass numbers, however. These are isotopes--- the same number of protons but a different number of neutrons • Examples: 1-H, 2-H, 3-H, C-12,13,14; U-235, 238 • Example: How many protons, neutrons, and electrons are in the following isotope of copper? Cu-63 Answer: 29 protons, 34 neutrons Chapter 2 Chem 1303 4 2.4. The Periodic Table • Figure 2.10 • organized by atomic number • General types of elements • metals, metalloids, nonmetals (locate) • main group elements (all 1A, 2A, 3 - 8 A) (locate) • Special names • 1A - alkali metal • 2A - alkaline earth metals • 7A - halogens • 8A - noble gases 2.5. Molecules and Ions 1. Molecule -- an aggregated of at least two atoms in a definite arrangement held together by chemical forces (bonds) • 7 diatomic elements: H, N, O, F, Cl, Br, I H2, N2, O2, F2, Cl2, Br2, I2 diatomic ball and stick model (Fig 2.12) space filling model (Fig 2.12) • Water molecule ("polyatomic") – draw 2. Ion - charged species formed from a neutral atom or molecule when electrons are gained or lost as a result of a chemical change • cations - positive • anions - negative • Examples Na and Cl # of protons and electrons in each • Common cations - Na+1, Ca+2, Al+3, Fe+2 or +3 • common anions: nonmetals: Cl-, O2Chapter 2 Chem 1303 5 2.6. Chemical Formulas - express the composition of molecules and ionic compounds in terms of chemical symbols • Levels of Structure • elemental composition • empirical formula • molecular formula • constitution • configuration • conformation • Molecular formula - shows the exact number of atoms of each element in the smallest unit of a substance Methane CH4 structural formula, ball and stick, space filling (Fig 2.12) Allotrope: one of two or more forms of an element, e.g., oxygen, O2 and ozone, O3 Empirical formula- tells which atoms are present and the simplest whole number ratio of the atoms (determined by experiment) glucose contains C, H, O in 1:2:1 ratio Empirical formula is CH2O other experiments show that molecular formula is six time the empirical form. i.e., C6H12O6 Chapter 2 Chem 1303 6 • Ionic compounds - discrete molecules are not present; ionic compounds are represented by their empirical formulas • Fig. 2.13 --- Structure of solid NaCl Ions are stacked in cube and touching each other • sum of positive charges equals the sum of negative charges • subscript of cation equals charge of anion and vice versa potassium bromide KBr one K+ and one Brzinc iodide ZnI2 aluminum oxide Al2O3 one Zn+2 and two Itwo Al+3 and three O-2 2.7. Naming compounds (Nomenclature) Four types of compounds: A. ionic B. molecular C. acids D. bases, hydrates Types of names • Common names - water is a common name: there is nothing in this name that tell us that it contains hydrogen and oxygen Formula Common name Systematic name AgCl Lunar caustic silver chloride H2SO4 MgSO4 Oil of vitriol sulfuric acid Epsom salts magnesium sulfate ⇒ Systematic names - sodium chloride is a systematic name; the name indicates what species are present Chapter 2 Chem 1303 7 Binary Compounds (two elements) IONIC Cmpds COVALENT Cmpds (molecular) (bonding- electrostatic) bonding - via shared electrons metals and nonmetals elements from left and right side of the periodic table nonmetals elements on same side (right side) of the periodic table ion charges show element ratios prefixes must be used to show element ratios + ion - name same as element - ion - drop ending & add ide Chapter 2 element on left named as element element on right - drop ending and add ide Chem 1303 8 General rules for naming binary compounds (2 elements) e.g., NaCl, FeBr3, Al2O3, N2O5, P4O10 A. Binary IONIC compounds : (One metal and one nonmetal) 2 elements- 2 words • 1st word metals named first - use name of element (no extra suffix) • 2nd word -nonmetal - use name of element plus ide suffix e.g., (remember that + charges = - charges to get a neutral complex) sodium bromide NaBr calcium oxide CaO barium chloride BaCl2 aluminum oxide Al2O3 Problem Write the formulas for potassium sulfide . • Common cations of metals use periodic table • +1 Li+1, Na+1, K+1, Ag+1 • +3 Mg+2, Ca+2,...Cu+2, Fe+2, Pb+2, Mn+2, Hg+2, Sn+2, Zn+2 Al+3, Cr+3, Fe+3 • others H+, NH4+, Hg2+ • +2 • Common anions of nonmetals • -1 F-1, Cl-1, Br-1, I-1 • -2 • -3 Chapter 2 O-2, S-2 N-3, P-3 Chem 1303 9 • Polyatomic Anions – two or more atoms combined in a single charged unit NH4+ H3O+ NO3PO43HCO3- ammonium ion hydronium ion nitrate ion phosphate ion hydrogen carbonate (or bicarbonate ion) LEARN THE 29 POLYATOMIC ANIONS (see Table at end of these notes, page 13) • But, some metals form more than one kind of ion. Then what? (usually a transition metal) • To name these types of metals, indicate the type of cation (charge) in parentheses using Roman numerals. e.g., MnO manganese (II) oxide Mn2O3 manganese (III) oxide MnO2 manganese (IV) oxide avoids problem with ous and ic • common metals with more than one type of cation Chapter 2 • +1, +2 Cu, Hg • +2, +3 Fe, Co, Ni • +2, +4 Sn, Pb • +2, +3, +6 Cr Chem 1303 10 Problem: Write the formulas for the following compounds: tin(II) fluoride SnF2 mercury (II) oxide HgO mercury (I) iodide Hg+ exists as Hg2+2 so Hg2I2 is the answer B. Binary MOLECULAR compounds : (Two nonmetals) • If 2 nonmetals are in different groups, the one to the left in the periodic table is named first • If 2 nonmetal are in the same group, the one farthest down that group in the periodic table is named first • The suffix -ide is added to the element named second • Counting prefixes (Table 2.4) are used with each name (Never use prefixes with IONIC compounds!!!!) • (mono is not used with the first name- it's "understood") Prefixes mono 1 hexa 6 di 2 hepta 7 tri 3 octa 8 tetra 4 nona 9 penta 5 deca 10 Examples Chapter 2 CO carbon monoxide CO2 SO2 carbon dioxide SO3 PCl3 sulfur trioxide PCl5 phosphorus pentachloride NO2 N2O4 nitrogen dioxide sulfur dioxide phosphorus trichloride dinitrogen tetroxide Chem 1303 11 Practice exercises NF3 Cl2O7 nitrogen trifluoride sulfur tetrafluoride SF4 dinitrogen pentoxide N 2 O5 dichlorine heptoxide C. Acids and Bases • An acid is a substance that yields hydrogen ions (H+) when dissolved in water 1. Some SIMPLE acids Table 2.5 Anion F- (fluoride) Corresponding Acid HF (hydrofluoric acid) Cl- (chloride) HCl (hydrochloric acid) Br- (bromide) HBr (hydrobromic acid) I- (iodide) HI (hydroiodic acid) CN- (cyanide) HCN (hydrocyanic acid) S2- (sulfide) H2S (hydrosulfuric acid) • Acids that contain hydrogen, oxygen and another element are called oxoacids. Chapter 2 Chem 1303 12 2. OXOACIDS Naming rules: • remove the ate and add ic • remove the ite and add ous • NO3- (nitrate) and NO2- (nitrite); • SO4-2(sulfate) and SO3-2 (sulfite), • other halogens follow chlorine example below Acetate ion acetic acid OCH3 OH CH3 C O O sulfite SO3-2 sulfurous acid H2SO3 or HOSOOH sulfate SO4-2 sulfuric acid H2SO4 HOSO2OH hypochlorous acid HOCl hypochlorite ClOchlorite ClO2- chlorous acid HClO2 HOClO chlorate ClO3- chloric acid HClO3 HOClO2 perchloric acid HClO4 HOClO3 perchlorate ClO4- Chapter 2 C Chem 1303 13 • A base is a substance that gives hydroxide anions (OH-) when dissolved in water. NaOH sodium hydroxide KOH potassium hydroxide Ba(OH)2 barium hydroxide NH3 ammonia NH3 + H2O D. NH4+ + OH- Hydrates • Hydrates are compounds that have a specific number of water molecules attached. BaCl2 . 2 H2O Sr(NO3)2 . 4 H2O barium chloride dihydrate strontium nitrate tetrahydrate 2.7. Introduction to Organic Compounds Hydrocarbons – compounds of carbon and hydrogen alkanes –Table 2.8 A. Alkanes CH4 C2H6 C3H8 C4H10 C5H12 Chapter 2 CnH2n+2 methane ethane propane butane pentane C6H14 C7H16 C8H18 C8H20 C10H22 hexane heptane octane nonane decane, etc...... Chem 1303 14 Selected Polyatomic Ions (Memorize these!) +1 -1 -2 NH4+ (ammonium) CN- (cyanide) O2-2 (peroxide) H3O+ (hydronium) OCN- (cyanate) C2O4-2 (oxalate) -3 SCN- (thiocyanate) OH- (hydroxide) CrO4-2 (chromate) C2H3O2- (acetate) Cr2O7-2 (dichromate) MnO4- (permanganate) NO3- (nitrate) NO2- (nitrite) H2PO4- (dihydrogen phosphate) HPO4-2 (hydrogen phosphate) PO4-3 (phosphate) HCO3- (hydrogen carbonate or bicarbonate) CO3-2 (carbonate) AsO4-3 (arsenate) HSO4- (hydrogen sulfate SO4-2 (sulfate) HSO3- (hydrogen sulfite) SO3-2 (sulfite) ClO4- (perchlorate) S2O3-2 (thiosulfate) ClO3- (chlorate) ClO2- (chlorite) ClO- (hypochlorite) + bromo, iodo relatives Chapter 2 Chem 1303 15 Periodic Table of the Elements 1 2 3 4 5 6 7 IA VIIIA (1) (18) 1 2 H IIA IIIA IVA VA VIA VIIA He 1.0080 (2) (13) (14) (15) (16) (17) 4.0026 3 4 5 6 7 8 9 10 Li Be B C N O F Ne 6.9410 9.0122 10.811 12.011 14.007 15.999 18.998 20.179 11 12 13 14 15 16 17 18 Na Mg IIIB IVB VB VIB VIIB . . VIIIB . . . IB IIB Al Si P S Cl Ar 22.990 24.305 (3) (4) (5) (6) (7) (8) (9) (10) (11) (12) 26.982 28.086 30.974 32.066 35.453 39.948 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 39.098 40.078 44.956 47.880 50.942 51.996 54.938 55.847 58.933 58.690 63.546 65.380 69.723 72.610 74.922 78.960 79.904 83.800 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 85.468 87.620 88.906 91.224 92.906 95.940 98.907 101.07 102.91 106.42 107.87 112.41 114.82 118.71 121.75 127.60 126.90 131.29 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 132.91 137.33 138.91 178.49 180.95 183.85 186.21 190.20 192.22 195.09 196.97 200.59 204.38 207.20 208.98 208.98 209.99 222.02 87 88 89 104 105 106 107 62 63 64 65 66 67 68 69 70 71 Fr Ra Ac Unq Unp Unh Uns 223.02 226.03 227.03 261.11 262.11 263.12 262.12 58 59 60 61 Chapter 2 Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 140.12 140.91 144.24 145.91 150.36 151.97 157.25 158.93 162.50 164.93 167.26 168.93 173.04 174.97 90 91 92 93 94 95 96 97 98 99 100 101 102 103 Th Pa U Np Pu 232.04 231.04 238.03 237.05 244.06 Am Cm 243.06 247.07 Bk Cf Es Fm Md No Lr 247.07 242.06 252.08 257.10 258.10 259.10 260.11 Chem 1303 16 Chapter 2 Chem 1303