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Molecules, Ions, and Their Compounds: Formulas, Names, and Properties

This chapter focuses on the interpretation, prediction, and writing of formulas for ionic and molecular compounds. It also covers naming compounds, understanding properties of ionic compounds, calculating molar mass and percent composition, and deriving formulas from experimental data.

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Molecules, Ions, and Their Compounds: Formulas, Names, and Properties

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  1. Chapter 3 Molecules, Ions, & their Compounds

  2. Chapter goals • Interpret, predict, and write formulas for ionic and molecular compounds. • Name compounds. • Understand some properties of ionic compounds. • Calculate and use molar mass. • Calculate percent composition for a compound and derive formulas from experimental data.

  3. Molecule • an assembly of 2 or more atoms (mostly of non-metals) bound together in a particular ratio and a particular manner • is the smallest identifiable unit into which a pure substance can be divided and still retain the composition and chemical properties of the substance

  4. Elements that exist as molecules • Atoms of most of the nonmetals form discrete molecules, except for the noble gases. • Some elements can exist in more than one form of molecule; the different forms are called allotropes. Examples are: • Diamond, graphite, and buckyballs for carbon. • O2 and O3 (ozone) for oxygen.

  5. ELEMENTS THAT EXIST AS MOLECULES Allotropes of C

  6. ELEMENTS THAT EXIST AS POLYATOMIC MOLECULES Sulfur: crown-shaped rings of S8 molecules White P4 and polymeric red phosphorus

  7. Molecular compoundsmodels - shapes Methane, CH4 Water, H2O Ammonia, NH3

  8. Molecular compounds • Are made of non-metals • H2O, carbon dioxide (CO2), ammonia (NH3), nitric acid (HNO3), ethanol (CH3CH2OH), sulfuric acid (H2SO4), glucose (C6H12O6), are examples among thousands • In the molecules of theses compounds, atoms share pairs of electrons

  9. Molecular compounds: Formulas NAME MOLECULAR CONDENSED STRUCTURAL FORMULA FORMULAFORMULA H H | | Ethanol C2H6O CH3CH2OHH─C ─ C─O─H | | H H H H | | Dimethyl C2H6O CH3OCH3H─C ─ O─ C─H ether| | H H Ethanol and dimethyl ether are said to be structural isomers.

  10. Ionic Compounds Ion • charged particle (atom or group of atoms) • cation: + charge • anion: – charge

  11. Ionic Compounds Sodium chloride or “table salt” is an example of an ionic compound.

  12. Ionic compounds • consist of positive and negative ions, mostly a metal and a non-metal, respectively. • have attractions called ionic bonds between positively (cations) and negatively charged ions (anions). • have high melting and boiling points. Tm of NaCl = 800 °C = 1472 °F • are solid at room temperature.

  13. Charge Balance Group IA VIIA NaCl, sodium chloride IIA of periodic table

  14. Periodic Table

  15. Monatomic Cations Metal atoms of group 1A lose one electron to produce a mono-positive ion

  16. Monatomic Cations Metal atoms of group 2A lose two electrons to produce a di-positive ion

  17. Monatomic Anions Nonmetals often gain one or more electrons and form ions having a negative charge equal to the group number of the element minus 8: Group Atom gained e- Resulting anion 5A N 3 N3- P 6A O 2 O2- S 7A F 1 F- Cl, Br, I Cl− , Br− , I−

  18. Charges of Representative Elements

  19. Monatomic Cations Transition metals (B-group elements) can form a no easily predictable variety of cations: Group Atom Electrons loss Resulting cation 7B Mn 2 Mn2+ 8B Fe 2 Fe2+ 8B Fe 3 Fe3+ 1B Cu 1 Cu+ 1B Cu 2 Cu2+ 2B Zn 2 Zn2+ 2B Cd 2 Cd2+

  20. Transition Metals form Positive Ions Most transition metals and Group 4A metals form 2 or more positive ions. Zn, Ag, and Cd form only one ion.

  21. Names of Some Common Ions Main group metals Nonmetal: change the last element name only part of name to ide

  22. Naming Cations • transition metals and In, Sn, Tl, Pb, Bi • new system: element name (charge in Roman numerals) • eg. Mn2+ manganese(II) • Mn3+ manganese(III) • Cr2+ chromium(II) Cr3+ chromium(III) • Fe2+ iron(II) Fe3+ iron(III) • Exceptions: when only one cation Ag+ silver Zn2+ zinc Cd2+ cadmium

  23. old system • Latin name-suffix • suffix = -ic for higher charge, -ous for lower charge • eg. Cu+ copper(I) cuprous ion • Cu2+ copper(II) cupric ion • Co2+ cobalt(II) cobaltous ion • Co3+ cobalt(III) cobaltic ion • Fe2+ iron(II) ferrous ion • Fe3+ iron(III) ferric ion

  24. Polyatomic Ions A polyatomic ion • is a group of atoms. • has an overall ionic charge, positive or negative.

  25. Polyatomic Ions (memorize) Some examples of polyatomic ions are NH4+ ammonium H3O+ hydronium OH− hydroxide N3−azide CO32− carbonate CN− cyanide CH3CO2− acetate C2O42− oxalate NO3−nitrate NO2− nitrite PO43− phosphate PO33−phosphite SO42− sulfate SO32−sulfite CrO42− chromate Cr2O72−dichromate MnO4− permanganate MnO42− manganate

  26. Hydrogenated Polyatomic Ions HCO3− hydrogen carbonate (bicarbonate) HSO4− hydrogen sulfate (bisulfate) HSO3− hydrogen sulfite (bisulfite) HPO42− hydrogen phosphate H2PO4− dihydrogen phosphate HS− hydrogen sulfide (from sulfide, S2−)

  27. Systematics • ClO4– perchlorate • ClO3– chlorate • ClO2– chlorite • ClO– hypochlorite The same with other halogens (except for F) • IO4– periodate • IO3– iodate • IO2– iodite • IO– hypoiodite

  28. PO53– perphosphate • PO43– phosphate • PO33– phosphite • PO23– hypophosphite • AsO53– perarsenate • AsO43– arsenate • AsO33– arsenite • AsO23– hypoarsenite

  29. Oxyanions that end in ate

  30. Other oxyanions:Once ions ending in ate are memorized, those with different number of O atoms are named with following prefixes and suffixes BrO– hypobromite BrO2– bromiteBrO3– bromateBrO4– perbromate

  31. Formula Unit • combination of ions in simplest whole number ratio to be electrically neutral (the simplest unit of an ionic compound) • consists of positively and negatively charged ions. • is neutral. • has charge balance. total positive charge = total negative charge The symbol of the metal is written first followed by the symbol of the nonmetal.

  32. Naming Ionic Compounds with Two Elements (binary compound) To name a compound that contains two elements, • identify thecationand anion. • name the cation first followed by the name of the anion.

  33. Examples, name each below • KBr • K+ • potassium • Br– • bromide • potassium bromide

  34. AlF3 • Al3+ • aluminum • F– • fluoride • aluminum fluoride

  35. Sr3P2 Sr P • Sr2+ 3(+2) = +6 2(−3) = −6 • strontium • P3– • phosphide • strontium phosphide

  36. CuCl2 Cu Cl • Cl– 1(+2) = +2 2(−1) = −2 • chloride • Cu2+ • copper(II) • cupric • copper(II) chloride • cupric chloride • CuCl: copper(I) chloride

  37. WF6 W F • F– w + 6x(−1) = 0 • fluoride w = 6+ • W6+ • tungsten(VI) • tungsten(VI) fluoride

  38. Naming Compounds with Polyatomic Ions The positive ion is named first followed by the name of the polyatomic ion. NaNO3sodiumnitrate K2SO4potassiumsulfate Fe(HCO3)3iron(III)bicarbonate or iron(III)hydrogen carbonate fe + 3x(−1) = 0 fe = 3+ iron(III) (NH4)3PO3ammoniumphosphite

  39. Writing Formulas with Polyatomic Ions The formula of an ionic compound • containing a polyatomic ion must have a charge balance that equals zero (0). Na+ and NO3− NaNO3 • with two or more polyatomic ions has the polyatomic ions in parentheses. Mg2+ and 2NO3− Mg(NO3)2 subscript 2 for charge balance • Aluminum sulfate 2Al3+ and 3SO42− Al2(SO4)3

  40. Learning Check Match each formula with the correct name. A. MgS 1) magnesium sulfite MgSO3 2) magnesium sulfate MgSO4 3) magnesium sulfide B. Ca(ClO3)2 1) calcium chlorate CaCl2 2) calcium chlorite Ca(ClO2)2 3) calcium chloride Name each of the following compounds: A. Mg(NO3)2 magnesium nitrate Mg2+ and 2 NO3− B. Cu(ClO3)2 copper(II) chlorate Cu2+ and 2 ClO3− C. PbO2 lead(IV) oxide Pb4+ and 2 O2− D. Fe2(SO4)3 iron(III) sulfate 2 Fe3+ and 3 SO42 − E. Ba3(PO3)2 barium phosphite 3 Ba2+ and 2 PO33 −

  41. Learning Check Select the correct formula for each. A. aluminum nitrate 1) AlNO3 2) Al(NO)3 3) Al(NO3)3 B. copper(II) nitrate 1) CuNO3 2) Cu(NO3)2 3) Cu2(NO3) C. iron(III) hydroxide 1) FeOH 2) Fe3OH 3) Fe(OH)3 D. tin(IV) hydroxide 1) Sn(OH)4 2) Sn(OH)2 3) Sn4(OH)

  42. CaSO4 • Ca2+ 1(+2) = +2 1(−2) = −2 • calcium • SO42– • sulfate • calcium sulfate

  43. (NH4)2S • NH4+ • ammonium • S2– • sulfide • ammonium sulfide

  44. (NH4)3PO4 • NH4+ • ammonium • PO43– • phosphate • ammonium phosphate

  45. Mo(BrO)6 • BrO– mo + 6x(−1) = 0 mo = 6+ • hypobromite • Mo6+ • molybdenum(VI) • molybdenum(VI) hypobromite

  46. Write formula for • rubidium bromide • Rb • Rb+ • Br • Br– • RbBr

  47. calcium phosphide • Ca • Ca2+ • P For neutrality we need • P3– 3 Ca2+ and 2 P3– • Ca3P2 that is, +6– 6 = 0

  48. niobium(IV) sulfite • Nb4+ In order to have the same • SO32–total + and – charge, we need • one Nb4+ ion and two SO32–, that is 1(+4) + 2 (–2) = 4 – 4 =0 (neutrality) • Nb(SO3)2 • barium phosphite: Ba2+ and PO33– for the compound to be neutral, we need 3 Ba2+ and 2 PO33–, that is, +6– 6 = 0 Ba3(PO3)2

  49. Electrostatic Forces The oppositely charged ions in ionic compounds are attracted to one another byELECTROSTATIC FORCES. These forces are governed byCOULOMB’S LAW.

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