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Unit 3 Language of Chemistry Part 1 Zumdahl: Chapter 4 Holt: Chapter 3 ATOMS: The Building Blocks of Matter Objectives 1. 2. 3. 4. 5. Law of conservation of mass Law of definite proportions Law of multiple proportions Dalton’s Atomic Theory How Dalton’s Atomic Theory relates to 1, 2, & 3 Atomic Theory Foundations Law of Conservation of Mass – mass is neither created or destroyed during a chemical or physical change Law of Definite Proportions – a compound contains the same proportions by mass regardless of the size of the sample Example: NaCl – always 39.34% Na & 60.66% Cl Atomic Theory Foundations Law of Multiple Proportions – if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers Example CO2 and CO : ratio of oxygen is always 2:1 JJ Thomson Showed atoms could emit negative particles “plum pudding” model Electrons were embedded in a positively charged spherical cloud Rutherford • Shoots alpha particles (Helium atoms) at gold foil • Expected to pass right through • Particles are deflected • Leads to idea of a dense positively charged center with e- orbiting around it Ernest Rutherford • Gold foil experiment Dalton’s Atomic theory 1. All matter is composed of atoms 2. Atoms of an element have the same size, mass and properties; atoms of a different element have different sizes, masses and properties 3. Atoms cannot be divided, created or destroyed 4. Atoms of different elements combine in simple whole number ratios 5. Chemical reactions combine, separate, or rearrange atoms Modern Atomic Theory • Atoms can be divided • Atoms of the same element can have different masses • All else remains the same Structure of the Atom Objectives Discovery of the Electron Rutherford’s Experiments Protons, Neutrons, Electrons Atomic Structure • Electron = no mass; negative charge • Proton mass = hydrogen atom; positive • Neutron mass = hydrogen atom; no charge • Dalton’s Model • JJ Thompson’s Plum Pudding Model The Electron • Mass of 9.109 x 10 -31 • Negative charge kg The Proton • Mass = 1.673 x 10 -27 • Positive charge kg The Neutron • Mass = 1.675 x 10 -27 • No charge kg Comparing Theories Dalton See notes Thompson Plum pudding model Rutherford Concept of the nucleus Electrons scattered thru positively charged cloud Positively charged Solid sphere Isotopes II.A.2(c) – compare the characteristics of isotopes of the same element IV.B.2(d) – calculate the weighted average atomic mass of an element from isotopic abundance, given the mass of each contributor Isotopes Def: atoms of the same element that have different masses Example: hydrogen protium – 1 proton in nucleus deuterium – 1 proton; 1 neutron tritium – 1 proton; 2 neutrons *Nuclide – general term for any isotope Writing Isotopes mass number Element Symbol atomic number Example – Uranium 235 235 92 U Average Mass Number Def: the weighted average of the atomic masses of the naturally occurring isotopes of an element Like calculating a “weighted” grade (decimal % of each isotope x mass of that isotope) Sample Calculation Isotope oxygen - 16 oxygen - 17 % abundance 99.762 0.038 Atomic Mass 15.994915 16.999131 oxygen - 18 0.200 17.999160 The Periodic Table IV.B.2(c) – use the periodic table to determine the atomic number, atomic mass, mass number, and number of protons, electrons and neutrons in isotopes of elements Atomic Number • Atomic Number – # of protons in nucleus • Element Symbol • Element Name • Atomic Weight • Electron Configuration 3 Li Lithium 6.941 [He]2s 1 Calculating Protons, Electrons & Neutrons IV.B.2(c) – use the periodic table to determine the atomic number, atomic mass, mass number, and number of protons, electrons and neutrons in isotopes of elements Mass Number Def: the number of protons and neutrons in the nucleus of an isotope mass # - atomic # = # of neutrons Example – oxygen Mass # (16) – atomic # (8) = # of neutrons (8) Formulas II.A.2(a) – use the IUPAC symbols of the most commonly referenced elements III.A.1(a) – distinguish between chemical symbols, empirical formulas, molecular formulas and structural formulas III.A.1(b) – interpret the information conveyed by chemical formulas for numbers of atoms of each element represented III.A.(d) – provide the interconversion of molecular formulas, structural formulas, and names, including common binary and ternary acids Chemical Symbols Molecular Formulas Empirical Formulas Structural Formulas Percent Composition III.A.1(c) – calculate the percent composition of a substance given its formula or masses of each component element in a sample % Composition Sample Problems Zumdahl – Ch 5 Holt – Ch 7 III.A.1(d) - provide the interconversion of molecular formulas, structural formulas, and names, including common binary and ternary acids Oxidation Numbers • • • • • • • Pure element have an oxidation number of zero Fluorine has an oxidation of -1 in all compounds Oxygen has an oxidation number of -2 except when in a peroxide when its oxidation number is -1 Hydrogen has an oxidation number of +1 except when bonded to a metal The algebraic sum of the oxidation numbers of all atoms in a neutral compound is zero The algebraic sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion These rules apply to covalently bonded atoms Oxidation Numbers Cont’d • • • • • • Group 1 = +1 Group 2 = +2 Transition Metals = varies Group 15 = varies (-3 usually) Group 16 = varies (-2 usually) Group 17 = varies (-1 usually) General Rules for Naming Chemical Compounds • Cation (pronounced cat-ion) – is named first – has a positive charge/oxidation number – name does NOT change • Anion (pronounced an-ion) – is named second – has a negative charge/oxidation number – ending is “-ide” Binary Ionic Compounds Type I • Metal is from Group 1 or 2 Examples: NaCl – sodium chloride KF – potassium fluoride CaBr2 – calcium bromide Li2O – lithium oxide Binary Ionic Compounds Type II • Metal is a transition metal • Many metals form more than one type of positive ion – – – – Iron: Fe2+ or Fe3+ Copper: Cu1+ or Cu2+ Lead: Pb2+ or Pb4+ Tin: Sn2+ or Sn4+ • Roman numeral indicates charge on the metal ion (NOT the number of atoms) • Look at 2nd element to determine charge on transition metal Examples of Type II EX: FeCl2 – iron(II) chloride FeCl3 – iron(III) chloride CuO – copper(II) oxide Cu2O – copper(I) oxide SnF4 – tin(IV) fluoride SnO2 – tin(IV) oxide SnO – tin(II) oxide **NOTE: sum of oxidation numbers = 0 Exceptions to the Rule Aluminum – always Al Cadmium – always Cd2+ Zinc – always Zn2+ Silver – always Ag+ 3+ No Roman numerals needed Binary Covalent Compounds or Binary Molecular Compounds • Formed between two non-metals 1.first element name doesn’t change 2.Second element ends with “-ide” 3.Don’t use “mono-” on first name 4.Use prefixes to indicate number of atoms of element Prefixes Mono– 1 Hepta- 7 Di– 2 Octa- 8 Tri- 3 Nona- 9 Tetra- 4 Deca- 10 Penta- 5 Undeca – 11 Hexa- 6 Dodeca - 12 Examples of Binary Covalent Compounds CO – carbon monoxide CO2 – carbon dioxide CCl4 – carbon tetrachloride SiO2 – silicon dioxide SeBr2 – selenium dibromide BF3 – boron trifluoride P2O4 – diphosphorous tetroxide P4O10 – tetraphosphorous decoxide Polyatomic Ions • Group of atoms with a shared charge III.A.1(c) – use the names, formulas, and charges of commonly referenced polyatomic ionsl Polyatomic Rule #1 • Common form ends with –ate Example: nitrate, chlorate, sulfate Polyatomic Rule #2 • If oxygen decreases by one, O-1, changes ending to “-ite” Example: nitrite, chlorite, sulfite Polyatomic Rule #3 • If oxygen decreases by two, O-2, add prefix “hypo-” • Keep ending “-ite” Example: hypochlorite Polyatomic Rule #4 • If oxygen increases by one, O+1, add prefix “per-” • Keep ending “-ate” Examples: perchlorate, permanganate Polyatomic Ions • • • • • • • • • • Chlorate Bromate Iodate Nitrate Permanganate Carbonate Silicate Selenate Phosphate Arsenate • • • • • • • • • • ClO3BrO3IO3NO3MnO4CO32SiO32SeO32PO43AsO43- More Polyatomic Ions • • • • • • • • • • Acetate Hydroxide Bicarbonate Bisulfate Ammonium Thiocyanate Thiosulfate Oxalate Chromate Dichromate C2H3O2OHHCO3HSO4NH4+ SCNS2O32C2O42CrO42Cr2O72- Binary Acids III.A.1(d) - provide the interconversion of molecular formulas, structural formulas, and names, including common binary and ternary acids Binary Acids • Hydrogen takes place of metal ion • Use prefix hydro- in place of hydrogen • Use suffix -ic on the root Examples – HCl – hydrochloric acid HF – hydroflouric acid H2S – hydrosulfuric acid Oxyacids III.A.1(d) - provide the interconversion of molecular formulas, structural formulas, and names, including common binary and ternary acids Oxyacids • contain hydrogen, a non-metal, and oxygen • usually contain 3 or 4 oxygen atoms • add suffix -ic to the stem Examples – HBrO3 – bromic acid H2CO3 – carbonic acid HClO3 – chloric acid HNO3 – nitric acid H3PO4 – phosphoric acid H2SO4 – sulfuric acid Oxyacids Cont’d • Common name ends with -ate • One extra oxygen add per- to stem name Example – HClO3 – chloric acid HClO4 – perchloric acid • One less oxygen change -ic ending to -ous Example – HClO3 – chloric acid HClO2 – chlorous acid Oxyacids cont’d • Two fewer oxygens change -ic ending to -ous AND add prefix hypoExamples HClO3 – chloric acid HClO – hypochlorous acid