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Dr. Lucio Gelmini Room 5-132A Gelminil@Macewan.Ca 497-5813 Chemistry 101 Dr. Lucio Gelmini Room 5-132A [email protected] 497-5813 http://academic.macewan.ca/gelminil Nuclear Atom Components of matter • Element – simplest type of substance with a unique identity (physical and chemical properties) – one type of atom – may be single atoms or molecules Components of matter • Compound – two or more different types of atoms – behave as a unit – chemically combined – unique physical and chemical identity • Mixture – two or more different types of elements and/or compounds physically intermingled Example of element vs. compound • Some properties of sodium, chlorine, and sodium chloride Elements combine by • Transferring electrons from an atom of one element to an atom of a different element TYPE OF COMPOUND FORMED? (ionic compound formation). • Sharing electrons between atoms TYPE OF COMPOUND FORMED? (covalent compound formation). Bonding Electrons involved (outer) not the nuclei in bonding Atom Identity • Atom identity is determined by #p+ in the nucleus (Z) • Isotope identity is determined by #n0 (A = Z + N) – Charge based on number of electrons e- Ion formation – ionic compounds • Atoms can gain or lose electrons – Lose: + ion = cation #p+ > #e- – Gain: − ion = anion #p+ > #e - • If an atom forms an ion… – Metals typically form cations (lose e-) – Nonmetals typically form anions (gain e-) Ion formation – ionic compounds Metals: lose #e- = A column number - Nonmetals: gain #e = 8 − A column number Example ions • What monatomic ions would the following elements most likely form? – sodium (column 1) → Na+ – chlorine (column 17) → Cl- – barium (column 2) → Ba2+ – nitrogen (column 15) → N3- Electron transfer • Electron transfer involves – loss and gain of e- – called oxidation and reduction, respectively (“redox”) Ionic bonding and ionic compounds NaCl = Na+Cl- • Oppositely charged ions: – attract each other – pair up in whole number ratios NaCl →zero overall charge – compounds are “ionic” • Examples: 2+ - CaF2 Ca F → 1(+2) + 2(−1) = 0 2+ 3- Ba3N2 Ba N → 3(+2) + 2(−3) = 0 + 2- K2OK O → 2(+1) + 1(−2) = 0 Differences between ionic and covalent bonding • Ionic substances are not molecular. – Exist as arrays of oppositely charge ions stabilized by ionic bonding (crystalline lattice) – No discrete units – ions are attracted to oppositely charged counterparts – “Formula unit” – lowest whole-number ratio that balances charges to zero • Covalent substances are molecular. – Discrete “units” of bonded atoms exist (“molecules”) H H C H H H C H H H Differences in compound formation Example compounds – row 3 combinations Covalent bonds • Covalent compounds form when atoms share e- • Usually among nonmetals • Attractive forces directional (between atoms) – Diatomic – 2 atoms O2, N2, CO,… – Tri atomic – 3 atoms O3, NO 2,H 2O… – Poly atomic – > 2 or 3 atoms CH4, H 2O2,C 4H 10,… – Diatomic and polyatomic ions exist - - + 2- OH , NO3 , NH4 , SO4 ,… − covalent bonds join atoms − group has gained or lost electrons Some nonmetal elements exist as molecules (1) (13)(2) (14) (15) (16)(18) (17) H2 N2 O2 F2 P4 S8 Cl2 Se8 Br2 I2 diatomic molecules mictetrato molecules ctatomico molecules Polyatomic ions • Polyatomic ions (learn them) – Contain covalently bonded atoms – Carry an overall charge – Act like ions (as a group) – Form ionic compounds • Example 2- – Carbonate ion, CO3 – Acts like ion with −2 charge – Calcium carbonate, CaCO3 2+ 2- cation = Ca ; anion = CO3 Compounds: formulas, names, masses • To effectively communicate in chemical terms, a chemical language is needed • Chemical symbols: “alphabet” • Specific combinations: “words” – Nickname = chemical formula – Full name = chemical name • Ionic and covalent compounds have different naming rules Binary Ionic (I) compound naming • Cation named first, followed by anion name – Cation same as the name of the parent metal • sodium metal → sodium ion Na → Na+ • calcium metal → calcium ion Ca → Ca2+ – Anion uses root of the nonmetal name and adds the suffix –ide • chlorine atom → chloride ion Cl → Cl− • oxygen atom → oxide ion O → O2− Compound examples (overall electrically neutral): NaCl sodium chloride CaO calcium oxide Na2Osodium oxide CaCl2 calcium chloride CN- cyanide OH- hydroxide Ionic (type II) naming with metals forming more than 1 ion • Hint: anion helps determine charge on metal overall, charges must add up to zero • Example. Name ions formula 4+ − tin(IV) fluoride Sn , F SnF4 + 2− copper(I) sulfide Cu , S Cu2S 3+ 2- iron(III) oxide Fe , O Fe2O3 Binary Ionic naming with metals forming more than 1 ion Example. Name ions formula 2+ − MnBr2 Mn , Br manganese(II) bromide 2+ 2 − PbS Pb , S lead(II) sulfide 2+ − CuI2 Cu , I copper(II) iodide Need to look at polyatomic ions To do that we need to know about naming acids Naming –Hydrates and Binary Covalent Compounds • Numerical (Greek) prefixes are used • Indicate number of each type of atom (or waters) Number Prefix Number Prefix Number Prefix 1 mono 4 tetra 8 octa 2 di 5 penta 9 nona 3 tri 6 hexa 10 deca 7 hepta • Examples: CO carbon monoxide CO2 carbon dioxide N2O dinitrogen monoxide P2O5 diphosphorous pentoxide SF6 sulfur hexafluoride IF5 iodine pentafluoride Naming Molecular Compounds • Similar to Binary ionic • More electropositive considered metal, and more electronegative atom = non metal • Various oxidation states, however, do not use roman numerals, use Greek prefix to identify number of each atoms • Only if there one “metal”atom do we drop the prefix, “non-metal”atoms always uses prefix Molecular compound naming • For example: N2O5 • More electropositive element named first, Nitrogen • Place Greek prefix in from of it (unless only one) Dinitrogen • More electronegative element named second, Oxygen • Place Greek prefix in from of it (unless only one) Pentoxygen • Finally to designate it is a binary compound, add “ide”ending Pentoxide OVERALL: Dinitrogen Pentoxide Hydrates contain bound water • copper(II) sulfate pentahydrate, CuSO4•5H2O Heating drives off waters blue CuSO4•5H2O → white anhydrous CuSO4 • Binary acids (Hn-X) – two groups to know Formula Pure In water Loss of H+ HF hydrogen fluoride Hydro fluoric F- fluoride ion acid HCl hydrogen chloride Hydro chloric Cl- chloride ion acid HBr hydrogen bromide Hydro bromic Br- bromide ion acid HI hydrogen iodide Hydro iodic l- iodide ion acid HCN hydrogen cyanide Hydro cyanic CN- cyanide ion acid 2- H2S hydrogen sulfide Hydro sulfuric S sulfide ion acid 2- H2Se hydrogen selenide Hydro selenic Se selenide ion acid 2- H2Te hydrogen telluride Hydro telluric Te telluride ion acid • OXY ACIDS (H-O-X) Look at X = halogen (group 17) X = S, Se or Te X = N or P,As X = C Organic acids (R-COOH) Other acids Oxoacids of Cl, Br and I Acid Name Ion(loss H+) Name - HClO4 perchloric acid ClO4 perchlorate anion - HClO3 chloric acid ClO3 chlorate anion - HClO2 chlorous acid ClO2 chlorite anion HClO hypochlorous acid ClO- hypochlorite anion NOTE: acidic hydrogens are generally listed 1st, not necessarily the way they bond (no H-Cl bond) Remember: Oxoacids must have a H-O bond HClO3 chloric acid = parent acid (often with three hydrogens) One less oxygen →“ous” ending acid. Hence, HClO2 chlorous acid “ic” ending acid→“ate” ending anion “ous” ending acid→“ite” ending anion Oxoacids of Cl, Br and I Acid Name Ion(loss H+) Name - HClO4 perchloric acid ClO4 perchlorate ion - HClO3 chloric acid ClO3 chlorate ion - HClO2 chlorous acid ClO2 chlorite acid HClO hypochlorous acid ClO- hypochlorite acid Acid Name Ion(loss H+) Name - HlO4 BrO4 - HBrO3 lO3 - HIO2 BrO2 HBrO IO- Oxoacids of Nitrogen • Only two: HNO3 and HNO2 parent acid higher oxidations state HNO3 has N(5+) and HNO2 has N(3+) - HNO3 Nitric acid → NO3 Nitrate anion - HNO2 Nitrous acid → NO2 Nitrite anion Oxoacids of P and As • Only two:H3PO4 and H3PO3 parent acid higher oxidations state H3PO4 3 has P(5+) and H3PO3 has P(3+) 3- H3PO4 Phosphoric acid → PO4 Phosphate anion 3- H3PO3 Phosphorous acid → PO3 Phosphite anion 3- H3AsO4 Arsenic acid → AsO4 Arsenate anion 3- H3AsO 3 Arsenous acid → AsO3 arsenite anion Oxoacids of P and As • Loss of only 1 or 2 hydrogen ions – List number of hydrogens - H2PO4 dihydrogen phosphate anion 2- HPO4 monohydrogen phosphate anion - H2PO3 dihydrogen phosphite anion 2- HPO3 monohydrogen phosphite anion - H2AsO4 dihydrogen arsenate anion 2- HAsO 4 monohydrogen arsenate anion - H2AsO3 dihydrogen arsenite anion 2- HAsO 3 monohydrogen arsenite anion Oxoacids of S, Se and Te • Only two:H2SO4 and H2SO3 parent acid higher oxidations state H2SO4 has S(6+) and H2SO3 has S(4+) 2- H2SO4 Sulfuric acid → SO4 Sulfate anion 2- H2SO 3 Sulfurous acid → SO3 Sulfite anion 2- H2SeO4 Selenic acid → SeO4 Selenate anion 2- H2SeO 3 Selenous acid → SeO3 Selenite anion 2- H2TeO4 Telluric acid → TeO4 Tellurate anion 2- H2TeO 3 Tellurous acid → TeO3 Tellurite anion Oxoacids of S, Se and Te • Loss of one hydrogen, may list number (mono) or omit - HSO4 (mono)hydrogen sulfate anion (bisulfate) - HSO3 (mono)hydrogen sulfite anion (bisulfite) - HSeO4 (mono)hydrogen selenate anion - HSeO3 (mono)hydrogen selenite anion - HTeO4 (mono)hydrogen tellurate anion - HTeO3 (mono)hydrogen tellurite anion Oxoacids of C • Only one important acid: H2CO3 = carbonic acid 2- Forms carbonate anion CO3 • Loss of one hydrogen, may list number (mono) or omit - HCO3 (mono)hydrogen carbonate anion (bicarbonate) Some consider carbonic acid as a hydrate of CO2 • Hence, H2CO3 is H2O CO2 and it is very unstable at room temperature breaking down → H2O(l) + CO2(g) Organic Acids • Sometimes difficult to tell • Organic acids have R –COOH Acetic acid is
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