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Unit 4 Formulas and Equations

Unit 4 Formulas and Equations

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Unit 4 Formulas and Equations

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  1. Unit 4Formulas and Equations Textbook Chapter 2, 6, & 8 Review Book Topic 2

  2. Chemical Symbols • Each element has been assigned a one-, two- or three- letter symbol for its identification • First letter is ALWAYS capitalized, additional letters are lowercase • Only recently discovered, unnamed elements are given three- letter symbols

  3. Some symbols show a relationship • Ex. Carbon ~ C Sodium ~ Na (Latin – natrium) • Symbols are assigned by IUPAC • International Union of Pure and Applied Chemists

  4. Roots used for naming elements: • 0 : nil • 1 : un • 2 : bi • 3: tri • 4 : quad • 5 : pent • 6 : hex • 7 : sept • 8 : oct • 9 : enn

  5. Ex. Element #109 • Un-nil-enn-ium (1)-(0)-(9) • Ex. Element #114

  6. Chemical Molecules • Monatomic molecules – uncombined elements, written without a subscript • Ex. Neon gas – Ne Argon gas – Ar • Diatomic molecules – elements can exist in nature as two identical atoms bonded together • Ex. Hydrogen – H2 (F, O, N, Cl, Br, I)

  7. Chemical Formulas • Chemists have identified over 10 million compounds • Compound – two or more elements that are chemically combined (bonded together) in definite proportions by mass • Ex. H2O, C6H12O6, H2O2

  8. No two compounds have identical properties

  9. Formulas use chemical symbols and numbers to show what elements and how many atoms of each are involved in each compound

  10. Chemical formula – shows the kinds and numbers of atoms in the smallest representative unit of the substance • If monatomic: use chemical symbol (ex. Kr) • If diatomic or a compound: use chemical symbols of elements involved, and subscripts to represent # of atoms present (ex. F2 or O3 or NaCl) • Types of formulas: molecular, empirical, structural

  11. Subscript – smaller number after an element symbol that indicates how many atoms of that element are in the molecule • Ex. H2O means there are 2 H and 1 O atom • Coefficient – number in front of a molecule’s formula indicating how many molecules are present • Ex. 2H2O means there are 2 water molecules

  12. Molecular formulas – shows the kinds and numbers of atoms present in a molecule of a compound • Subscript written after the symbol indicates the # of atoms of each element • If only 1 atom, subscript of 1 is omitted • Show composition but NOT molecular structure

  13. Empirical formula (“formula unit”) – shows the lowest whole number ratio of ions in a compound • Ex. MgCl2 • For every 1 Mg+, there are 2 Cl- • Ex. H2O and H4O2 • Both have a ratio of 2 H : 1 O

  14. Molecular formulas can be seen as a multiple of an empirical formula • Ex. Glucose: C6H12O6 (molecular) CH2O (empirical) 6(CH20) = C6H12O6

  15. Structural formula – shows the physical organization of the atoms in a molecule

  16. Law of definite proportions – in any compound, the masses of the elements involved are always in the same proportions • Ex. NaCl always has 1 Na (23 amu) and 1 Cl (35 amu) = 58 amu total for one NaCl

  17. Ex. H2O always has 2 H (total 2 amu) and 1 O (16 amu) = 18 amu total for one H2O • Proportions of mass equals the ratio proportions of the number of atoms of each element in the molecule

  18. Law of multiple proportions – whenever two elements form more than one compound (ex. H2O and H2O2), the different masses of one element (ex. O versus O2) that combine with the same mass of the other element (H2) are in the ratio of small whole numbers • Ex. We have two compounds, each with 2 g of element B. Compound 1 has 5g element A, compound 2 has 10 g element A

  19. Review Questions • Determine the empirical formula for the compounds below: • C6H8O6 (vitamin c) • Hg2Br2 • K2CrO4 • N2H4 • C6H6 • C2H4O2

  20. Review Questions • Determine the empirical formula for the compounds below: • C6H8O6 (vitamin c) C3H4O3 • Hg2Br2 HgBr • K2CrO4 K2CrO4 • N2H4 NH2 • C6H6 CH • C2H4O2 CH2O

  21. Atoms, Compounds and Ions • Atoms and compounds are electrically neutral • (# p+ = # e-)

  22. Ions have a net charge, either (+) or (-) • (# p+≠ # e-) • (+) ions attract (-) ions in a ratio that produces a neutral compound

  23. Monatomic Ions • Ions consisting of only one atom • Ionic charges are found using the periodic table (look at group #s) • Metals have a (+) ionic charge, nonmetals have a (-) ionic charge

  24. Metallic elements tend to lose electrons (forming cations) • Group 1: 1+ charge • Group 2: 2+ charge • Aluminum: 3+ charge

  25. Nonmetallic elements tend to gain electrons (forming anions) • N, P, and As: 3- charge • O, S and Se: 2- charge • F, Cl, Br, I (group 17): 1- charge • Nonmetal ionic charge is found by subtracting the group number (in the form of 5A, 6A, 7A, etc.) from 8

  26. Group 0 usually does not form ions (noble gases) • Transition metals tend to have more than one ionic charge (represented by the oxidation numbers on the periodic table)

  27. Determining the ionic charge of transition metals: • Roman numerals are used in parentheses to indicate the numerical charge • Form: Element name(roman numeral) ion • No spaces are used between the element name and the first parentheses • Ex. Cu 2+: Copper(II) ion Sn4+: Tin(IV) ion

  28. Roman Numerals:

  29. Polyatomic Ions • Tightly bound group of atoms that behave as a unit • Carry an overall charge (+ or -) • Reference Table E

  30. Names usually end in –ite or –ate • Three exceptions: • Ammonium cation (NH4+) • Two polyatomic anions ending in –ide • Cyanide (CN-) • Hydroxide ion (OH-)

  31. -ite/-ate pairs of polyatomic ions: -ite-ate SO32-, sulfite SO42-, sulfate NO2-, nitrite NO3-, nitrate ClO2-, chlorite ClO3-, chlorate • The –ite ending indicates one less oxygen atom than the –ate ending

  32. When the formula for a polyatomic ion begins with a hydrogen ion (H+): • The charge on the new ion is the sum of the ionic charges: • H+ + CO32- HCO3- (hydrogen carbonate) • H+ + PO43- HPO42- (hydrogen phosphate) • H+ + HPO42- H2PO4- (dihydrogen phosphate)

  33. Review Questions • How can the periodic table be used to determine the charge of an ion? • Explain what is meant by a polyatomic ion

  34. Review Questions • How can the periodic table be used to determine the charge of an ion? Look up the oxidation #s to see the different ion possibilities for each element or look at the group number • Explain what is meant by a polyatomic ion

  35. Review Questions • How can the periodic table be used to determine the charge of an ion? Look up the oxidation #s to see the different ion possibilities for each element or look at the group number • Explain what is meant by a polyatomic ion Contains more than one ion but acts as a single “unit” by carrying an overall charge

  36. Using only your periodic table, write the formula for the typical ion of each element and identify it as an anion or cation: • Potassium • Sulfur • Argon • Bromine • Beryllium • Sodium

  37. Using only your periodic table, write the formula for the typical ion of each element and identify it as an anion or cation: • Potassium K+ cation • Sulfur S2- anion • Argon NONE • Bromine Br1- anion • Beryllium Be2+ cation • Sodium Na+ cation

  38. Write the formula, including the charge, for each ion: • Ammonium ion • Tin(II) ion • Chromate • Nitrate ion • Cyanide ion • Iron(III) ion • Permanganate ion • Manganese(II) ion

  39. Write the formula, including the charge, for each ion: • Ammonium ion NH4+ • Tin(II) ion Sn2+ • Chromate CrO42- • Nitrate ion NO3- • Cyanide ion CN- • Iron(III) ion Fe3+ • Permanganate ion MnO4- • Manganese(II) ion Mn2+

  40. Ionic Compounds • Monatomic or polyatomic ions attract each other in a ratio that produces a neutral compound • Opposite charges attract ! • Ex. HCl (H+ + Cl-) • Ex. H2SO4(H22+ + SO42-) • Ex. AgNO3 (Ag+ + NO3-)

  41. A compound’s name should indicate its composition, behavior, and how it is related to other compounds….common names do not tell us anything about chemical composition! • Ex. Sodium chloride versus salt • Ex. Dihydrogen oxide versus water

  42. Binary compounds – composed of two elements • (+) charge of the cation must balance with the (-) charge of the anion (equal but opposite charges) • Net ionic charge = 0

  43. Formulas for ionic compounds are usually written with the cation first, followed by the anion and ALWAYS show the lowest whole-number ration of ions in the compound

  44. Ternary ionic compounds – composed of atoms of three different elements • Usually contains a polyatomic ion • Parentheses can be used around the polyatomic ion to show if more than one are used in a reaction

  45. Writing Ionic Compound Formulas • Crisscross method: • The charge of each ion is crossed over and used as a subscript for the other ion • For many elements, the oxidation state is equal to the charge on the ion • The signs (+ or -) are dropped when used as subscripts

  46. Review Questions • Create a molecular formula between the following ions: • Fe3+ O2- • Ca2+ S2- • Ba2+ S2- • Li+ O2- • Ca2+ N3- • Cu2+ I- • K+N3-

  47. Review Questions • Create a molecular formula between the following ions: • Fe3+ O2- Fe2O3 • Ca2+ S2-Ca2S2 = CaS • Ba2+ S2-Ba2S2 = BaS • Li+ O2-Li2O • Ca2+ N3-Ca3N2 • Cu2+ I- CuI2 • K+N3- K3N

  48. Ex. Na+Cl- • Ex. Mg2+Cl- • Ex. Ca2+ NO3- • Ex. Na+ NO3-

  49. Ex. Na+Cl- NaCl • Ex. Mg2+Cl- MgCl2 • Ex. Ca2+ NO3- Ca(NO3)2 • Ex. Na+ NO3- NaNO3