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Chemical Formulas and Naming

Chemical Formulas and Naming. Unit 4. Ions. Atoms that have a positive or negative charge To become an ion, an atom gains or loses electrons Cation = positively charged ion; Lose electron (s) Metals form positive ions Anion = negatively charged ion; Gain electron (s). Oxidation Numbers.

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Chemical Formulas and Naming

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  1. Chemical Formulas and Naming Unit 4

  2. Ions • Atoms that have a positive or negative charge • To become an ion, an atom gains or loses electrons • Cation = positively charged ion; • Lose electron (s) • Metals form positive ions • Anion = negatively charged ion; • Gain electron (s)

  3. Oxidation Numbers • Indicate the charge on the ion • Found on the periodic table • Common oxidation numbers are given in additional table • Note: • Group 8 elements do not form ions; No ox # • Transition metals have multiple ox #’s

  4. Polyatomic Ions • Def: tightly bound groups of atoms that behave as a unit and carry a charge • Ion composed of more than one atom • List of common polyatomic ions

  5. Polyatomic Ions +1 • NH4+ Ammonium -1 • NO3- Nitrate • NO2- Nitrite • CN- Cyanide • OH- Hydroxide • CH3COO- or C2H3O2- Acetate • MnO4- Permanganate • ClO3-Chlorate -2 • CO3-2 Carbonate • SO4-2 Sulfate • C2O4-2 Oxalate • CrO4-2 Chromate • Cr2O7–2 Dichromate -3 • PO4-3 Phosphate

  6. Subscript “2” Indicates how many atoms of the preceding element are present. The coefficient of “3” indicates the number of molecules Compounds • What is a compound? • Compounds are composed of elements. • They are written by using chemical symbols and subscripts. • Chemical symbols are a capital letter or a capital letter and a lower case letter, that stands for an element. 3CO2 • Chemical Symbols “C” and “O”

  7. Compounds • How many atoms are present in the following examples? CH4 One carbon and four hydrogen = 5 atoms C6H12O6 Six carbon, twelve hydrogen and six oxygen = 24 atoms total NaOH One sodium, one oxygen and one hydrogen = 3 atoms total

  8. Chemical Formulas • The elements are listed from most metallic on the left, to least metallic on the right. • Example; For a compound made from sodium and oxygen. Sodium is listed first and oxygen second. Na2O For a compound made from nitrogen and calcium. Calcium is listed first and nitrogen second. Ca3N2

  9. Writing Formulas • Compounds are electrically neutral • Positive ion charge must balance the negative ion charge • Ex: Write the formula for a compound made from: • potassium and chlorine • calcium and bromine

  10. Criss Cross Method • Ion charge crosses over to use as the subscript on the opposite atom • Useful when charges don’t balance out easily • Ex: aluminum and oxygen • Al2O3

  11. Chemical Formula Problems • What are the formulas for compounds made from the following elements? K3N Potassium and nitrogen Na2O Sodium and oxygen Rb3P Phosphorus and rubidium Al2S3 Aluminum and sulfur

  12. Molecular and Ionic Compounds • Ionic compounds are made from metal (cation)/ non-metal (anion) combinations. • Molecular compounds (Covalent compounds) are made from non-metal/ non-metal combinations.

  13. Naming Compounds: IUPAC Nomenclature • International Union of Pure and Applied Chemistry • Systematic approach to naming compounds • Same names everywhere • Common names don’t tell us about their chemical composition • Ex: baking soda, quicksilver, lye, laughing gas

  14. Naming Compounds: Binary Ionic • Compound with only 2 different ions • 1 metal and 1 nonmetal • 1 positive ion and 1 negative ion • Name the cation and anion separately • Change the ending of the non-metal to “ide” • Example; • A compound made from chlorine and lithium becomes Lithium chloride

  15. Name the Following Compounds. K3N Potassium nitride Na2O Sodium oxide Rb3P Rubidium phosphide Al2S3 Aluminum sulfide

  16. Naming with Transition Metals (Binary Ionic) • Most metallic element is listed first. • If cation has more than one oxidation #, use Roman Numerals to indicate the charge • Ex: Fe = iron, Fe+2 = iron (II), Fe+3 = iron (III) • Anion = name of the element ending in “-ide” • Ex: O = oxygen; O-2 = oxide • Ex: F = fluorine; F-1 = fluoride

  17. Examples with Transition Metals FeCl2 Iron (II) chloride Cobalt (ll) carbonate CoCO3 Manganese(Vl) oxide Mn2O6 Strontium chloride SrCl2

  18. Naming Compounds: Tertiary Ionic • Compounds with polyatomic ions • Separate cation from anion and name individually • Most metallic element is listed first. • The ending of the polyatomic remains the same. • Name of cation = name of element • Name of anion = name of polyatomic ion • Ex: CaCO3 Mg(NO3)2 Na2SO3

  19. Naming Tertiary Compounds • Examples Potassium nitrate KNO3 Sodium oxalate Na2C2O4 Rubidium phosphate Rb3PO4 Ammonium sulfate (NH4)2SO4

  20. Naming Compounds: Binary Molecular • Compound with only 2 different atoms • 2 nonmetals • Can combine in more than one way with the same elements • Ex: CO and CO2 • Name the same way as ionic compounds (Most metallic element is listed first) • Need to include Greek prefixes to designate how many atoms of each element are present

  21. Greek Prefixes for Binary Molecular Compounds Mono = 1 Di = 2 Tri = 3 Tetra =4 Penta = 5 Hexa = 6 Hepta= 7 Octa = 8 Nona=9 Deca = 10

  22. Examples of Molecular Compounds P2O5 diphosphorus pentoxide N2O dinitrogen monoxide monocarbon dioxide or carbon dioxide CO2 CO carbon monoxide

  23. Summary of Naming And Formula Writing • In an ionic compound, the ionic charges add to zero. • An “-ide” ending indicates a binary compound. • An “-ite” or “-ate” ending indicates a polyatomic anion is in the formula. • Roman numerals are used to indicate the charges on transition metals. • Greek prefixes are used to indicate the number of atoms present in a molecular compound.

  24. Balancing Equations: WHY? • Law of Conservation of Matter – matter can neither be created nor destroyed  • Anything on the left hand side must be found on the right hand side – cannot magically lose or gain atoms

  25. Balancing Equations: Coefficients • What is a coefficient? • Number found in front of a compound’s formula • Indicates how many molecules are involved in the reaction • Add coefficients to balance the equation • Versus subscripts: • subscripts indicate how many atoms in a molecule (ex: CaBr2 – 2 Br atoms) • coefficients indicate how many molecules in a reaction (ex: 2CaBr2 – 2 CaBr2 molecules; 4 Br atoms)

  26. Symbols Used in Equations + - separates 2 reactants or 2 products  - separates reactants from products; “yields” (s) - solid state (l) - liquid state

  27. Symbols Used in Equations cont. (aq) - aqueous solution (g) - gaseous state  - heat Formula written above or below yield sign - catalyst

  28. Balancing Equations • Example HgO (s) Hg (l) + O2 (g)

  29. Balancing Equation Examples • __C4H10 + __O2 __CO2 + __H2O • For these equations balance H then C then O (known as combustion equations) __CaCO3 + __HCl __ CO2 + __CaCl2 + __H2O

  30. Types of Reactions • There are millions of different chemical reactions • We need to learn how to recognize patterns of chemical behavior to predict products in many reactions • Most reactions can be divided into five main types: 1- Synthesis (or Combination) 2- Decomposition 3- Single displacement reactions 4- Double displacement reaction 5- Combustion

  31. Synthesis (Combination) Reactions • In a synthesis reaction, two or more substances combine to form another substances (reactant + reactant  1 product) • General Formula: A + B AB • Example: 2H2 (g) + O2 (g) 2H2O

  32. Synthesis (Combination) Exceptions • Reaction of metal oxide and water yields a metal hydroxide (OH-1) • Examples MgO + H2O Mg(OH)2 CaO + H2O Ca(OH)2 LiO + H2O Li(OH)

  33. Decomposition Reactions • Occurs when one substance breaks down or decomposes into two or more substance (1 Reactant  Product + Product) • General Formula: AB A + B • Example: 2 H2O (l) 2H2 (g) + O2 (g)

  34. Decomposition Exceptions Carbonates, chlorates, and hydroxides are special case decomposition reactions that do not go to the elements. • Carbonates (CO32-) decompose to carbon dioxide and a metal oxide • Example: CaCO3  CO2 + CaO • Chlorates (ClO3-) decompose to oxygen gas and a metal chloride • Example: 2 Al(ClO3)3  2 AlCl3 + 9 O2 • Hydroxides (OH-1) decompose to metal oxide and water • Example: Ca(OH)2 CaO + H2O

  35. Single Replacement Reactions • A single displacement reaction occurs when one element replaces another element in a compound. • General Formulas • 1st case: A + BC AC + B (A is a metal) • 2nd case: A + BC BA + C (A is a nonmetal) • Example: Cu(s) + 2AgNO3(aq) Cu(NO3)2 (aq) + 2Ag(s)

  36. Single Replacement Exceptions • Reaction of a metal and an acid produces an ionic compound and hydrogen gas (H2) • Example Mg + HCl  MgCl2 + H2 • Reaction of a metal with water produces a metal hydroxide (OH-1) and hydrogen gas (H2) • Example Ca + H2O  Ca(OH)2 + H2

  37. Double Replacement Reactions • A double displacement reaction takes place when two ionic compounds in solution combine • In this type of reaction a precipitate forms • This type of reaction yields a solid (an insoluble compound that doesn’t dissolve) and a new solution (soluble-dissolves in water) • How do we determine if something is soluble or insoluble? USE YOUR SOLUBILITY RULES (TABLE) • More about this later…. • The positive ion of one compound replaces the positive ion of the other to form two new compounds. • Compound + compound  compound+ compound • General Formula: AB + CD  AD + CB

  38. Double Replacement Reactions • Think about it like “foil”ing in algebra, first and last ions go together + inside ions go together • Example: AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq) • Another example: K2SO4(aq) + Ba(NO3)2(aq)  KNO3(aq) + BaSO4(s) 2

  39. Combustion Reactions • Combustion reactions occur when a hydrocarbon (an organic compound) reacts with oxygen gas to produce carbon dioxide gas (CO2) and water (H2O) • Used to heat homes and run automobiles (octane, as in gasoline, is C8H18) • General Formula: hydrocarbon + oxygen ----> carbon dioxide and water CxHy + O2  CO2 + H2O • Example (Start balancing with hydrogen, then carbons, and finally oxygen) CH4(g) + 2O2(g) ----> 2H2O(g) + CO2(g) 2C2H6(g) + 7O2(g) ----> 6H20(g) + 4CO2(g)

  40. Edgar Allen Poe’s drooping eyes and mouth are potential signs of CO poisoning.

  41. Acid Nomenclature Binary  Ternary An easy way to remember which goes with which… “In the cafeteria, youATEsomethingICky”

  42. Naming Acid Examples hydrobromic acid • HBr (aq) • H2CO3 • H2SO3  carbonic acid  sulfurous acid

  43. Organic Compounds (Hydrocarbons) • Contain carbon surrounded by hydrogen's • Have covalent bonds • Have low melting points • Have low boiling points • Burn in air (oxygen) • Are soluble in nonpolar solvents • Form large molecules

  44. Organic Compounds (Hydrocarbons) • Contain C and H only • Contain single bonds C-C • Have 4 bonds to every carbon (C) atom • Are nonpolar

  45. Complete Structural Formulas Show the bonds between each of the atoms H H   H  C  H H C H   H H CH4 , methane

  46. IUPAC Names Name # carbons Structural Formula Methane 1 CH4 Ethane 2 CH3CH3 Propane 3 CH3CH2CH3 Butane 4 CH3CH2CH2CH3 Pentane 5 CH3CH2CH2CH2CH3

  47. IUPAC NAMES Name # carbons Structural Formula Hexane 6 CH3CH2CH2CH2CH2CH3 Heptane 7 CH3CH2CH2CH2CH2CH2CH3 Octane 8 CH3CH2CH2CH2CH2CH2CH2CH3 Nonane 9 CH3 CH2 CH2CH2CH2CH2CH2CH2CH3 Decane 10 CH3CH2CH2CH2CH2CH2CH2CH2CH2CH3

  48. IUPAC NAMES • Sometimes there’s a branch off of a main chain • To name a branch use the same prefixes on the previous 2 slides, but the suffix becomes –yl • For example: • 1 carbon branch would be methyl • 2 carbon branch would be ethyl

  49. Mnemonic for First Four Prefixes First four prefixes Meth- Eth- Prop- But- • Monkeys • Eat • Peeled • Bananas

  50. General Formulas • Alkanes = CnH2n+2(formula for hydrocarbons) • Alkenes = CnH2n • Alkynes = CnH2n-2

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