Honors Chemistry - Topic IV Part I Study Guide The Language of Chemistry I. Chemical Nomenclature A. Inorganic Nomenclature (pp.11132; 349-350, 554-555; 213-220) 1. Types of formulas chart showing positions of these ions. a. binary ionic compounds type I and type II Type I (1): Representative Metal and Nonmetal tend to lose valence electrons, so positively charged ion (+) called a cation Metallic ions from groups 1 and 2 on the Periodic Table. They have only one type of charge. They binary ionic compounds formed are composed of a positive ion (cation) that is written first and a negative ion (anion). The following rules show how to name a write the formulas for binary ionic compounds. CaCl2 is used as an example. 1. Name the cation first and then the anion. 2. The monoatomic (one-atom) cation takes its name from the name of the element. Therefore the calcium ion, Ca 2+, is called calcium and its chemical symbol appears first. 3. The monoatomic anion with which the cation combines is named by taking the root of the elements name and adding ide. You must know this rule. The anion s name comes second. Therefore, the chlorine ion Cl, is called chloride. 4. The name of this compound is calcium chloride. A quick way to determine the formula of a binary ionic compound is to use the crisscross rule. 1. Take the numerical value of the metal ion s superscript (forget about the charge symbol) and move it to the bottom right-hand side of the nonmetal s symbol as a subscript. 2. Then take the numerical value of the nonmetal s superscript and make it the subscript of the metal.
o Make sure you reduce al subscripts by a common factor. For example Mg 2+ O Mg2O2 simplifies to MgO.) Examples of Category I Binary Ionic Compounds: Ions Present Formula Name K +, Cl - KCl Potassium chloride Na +, I - NaI Sodium iodide Ca 2+, S CaS Calcium sulfide Al 3+, F - AIF3 Aluminum Fluoride Li +, N 3- Li3N Lithium Nitride Type Ii (2): Transition Metal and nonmetal In category II binary ionic compounds, the metals form more than one ion, each with a different charge. The metallic ions (cation) ionically bind with the negatively charged ion (anion). The table bellows lists most of the metals that form more than one type of ionic cation and therefore more than one binary ionic compound with a given anion. Ion Systematic Name Fe 3+ Iron (III) Fe 2+ Iron (II) Cu 3+ Copper (III) Cu 2+ Copper (II) Sn 4+ Tin (IV) Sn 2+ Tin (II) Pb 4+ Lead (IV) Pb 2+ Lead (II) Hg2 2+ Mercury (I)* Hg 2+ Mercury (II) * this form of mercury (I) ions always occurs bonded together. Although these metals are transition metals, they form only one of cation. So a Roman numeral is not used when naming their compounds. (you see this on periodic table as well) o Ag 1+ Silver o Zn 2+ Zinc o Cd 2+ Cadmium The names of these ions (besides the 3 above) have Roman Numerals in parenthesis to indicate the charge of the metallic ion used as the cation.
Examples of Category II Binary Ionic Compounds: Formula Name CuCl Copper (I) chloride HgO Merucry (II) oxide FeO Iron (II) oxide MnO2 Manganese (IV) oxide PbCl2 Lead (II) chloride b. binary molecular (covalent) compounds 2 nonmetals Don t contain ions To name binary covalent compounds, use these steps 1. The 1 st element In formula is prefix (if applicable) and full elemental name. 2. The second element uses prefix and its elemental name and ends in ide. 3. Prefixes are used to detonate the number of the element present. Those prefixes are shown below. 4. You never use mono if the 1 st one is singular. For example CO is called carbon monoxide. o For ease of pronunciation, the final a or o of the prefix is dropped if the element begins with a vowel. (Drop double vowels-see example below) o To write the formula, use the same steps used for ionic compounds Examples: o CO Carbon monoxide 1 oxygen o CO2 Carbon dioxide 2 oxygen o NO Nitrogen monoxide o N2C4 Dinitrogen tetroxide (drop double vowel) o N2O Dinitrogen monoxide o NO2 Nitrogen dioxide Prefixes 1. mono 2. di 3. tri 4. tetra 5. penta 6. hexa 7. hepta 8. octa 9. nona 10. deca
c. ionic compounds with polyatomic ions (see chart of polyatomic ions below 1+ CHARGE CHARGE CHARGE 3- CHARGE ion name ion name ion name ion name 1+ NH 4 ammonium HCO 3 hydrogen carbonate CO 3 carbonate 3- PO 3 phosphite H 3 O 1+ hydronium HSO 3 hydrogen sulfite SO 3 sulfite 3- PO 4 phosphate 2+ Hg 2 mercury(i) HSO 4 hydrogen sulfate SO 4 sulfate NO 2 nitrite NO 3 nitrate C 2 O 4 oxalate OH hydroxide CrO 4 chromate C 2 H 3 O 2 acetate Cr 2 O 7 dichromate CrO 2 chromite CN cyanide O 2 peroxide MnO 4 ClO ClO 2 ClO 3 ClO 4 permanganate hypochlorite chlorite chlorate perchlorate IO 3 iodate N 3 azide 2. More Nomenclature a. common name v. systematic name Why use naming rules at all? o common names do not discuss chemical composition o systematic names are a standardized way of naming compounds Some common names and their chemical (systematic) names: asbestos = magnesium silicate table salt = sodium chloride aspirin = acetylsalicylic acid quicksilver = mercury baking soda = sodium bicarbonate silica = silicon dioxide black lead = graphite soda ash, dry = dry sodium carbonate borax = sodium borate soda lye = sodium hydroxide brine = strong sodium chloride solution soluble glass = sodium silicate chalk = calcium carbonate talc or talcum = magnesium silicate drinking alcohol = ethanol vinegar = dilute acetic acid Epsom salts = magnesium sulfate Vitamin C = ascorbic acid laughing gas = dinitrogen monoxide water = H2O lime = calcium oxide water glass = sodium silicate lime, slaked = calcium hydroxide limewater = calcium hydroxide solution
b. stock system v. classical system page 118 Lower Charge Stock System Higher Charge Element Formula Name Formula Name #$%%&'((((((((((((((((#) + *$%%&'(+,- #) 2+ *$%%&'(+,,-,'$. /& 2+ 0'$.+,,- /& 3+ 0'$.+,,,- 1&23( 45 2+ 6&23(+,,- 45 4+ 6&23+,7-8&'*)'9 :; < <= >&'*)'9+,- :; 2+ >&'*)'9+,,-?0. @. 2+?0.+,,- @. 4+?0.(+,7- "! Ion Names: Classical System Lower Charge Higher Charge Element Formula Name Formula Name #$%%&' #) = *)%'$)F #) <= #)%'0*,'$. /& <= H&''$)F /& D= H&''0* 1&23 45 <= %6)>5$)F 45 "= %6)>50* 8&'*)'9 :; >&'*)'$)F :; <= >&'*)'0*?0. @. <= FG2..$)F @. "= FG2..0*,.(GI&(#62FF0*26(@9FG&>(GI&(.2>&($H(GI&(>&G26(+)F)2669(GI&(12G0.(.2>&-((0F(>$30H0&3(J0GI(GI&(F)HH0K&F(-ous 2.3(icL AE 8&G26(.2>&(&.3F(0. -ous 6$J&'(*I2';& -ic I0;I&'(*I2';&.$.>&G26(.2>&(&.3F(0. -ide c. naming acids and bases Names and formulas for common acids Acids are substances, which produce H+ ions in water. We name the acids based on the negative ion which is left over after the H+ ion is given off.
Rules Negative ion suffix ate ite Acid ic acid ous acid ide Rules presented in a different way. Rule -ic acids form ate salts. -ous acids form ite salts. hyro-(stem)-ic acids form ide salts. Formulas hydro ic acid Example Sulfuric acid forms sulfate salts. Sulfurous acid forms sulfite salts. Hydrochloric acid forms chloride salts. (HCl) Names H2SO4 (aq) (sulfate ion) Sulfuric acid H2SO3 (aq) (sulfite ion) Sulfurous acid HNO3 (aq) (nitrate ion) Nitric acid HNO2 (aq) (nitrite ion) Nitrous acid H3PO4 (aq) (phosphate ion) Phosphoric acid HClO4 (aq) (perchlorate ion) Perchloric acid HClO3 (aq) (chlorate ion) Chloric acid HClO2 (aq) (chlorite ion) Chlorous acid HClO (aq) (hypochlorite ion) Hypochlorous acid HCl (aq) (chloride ion) Hydrochloric acid HF (aq) (fluoride ion) Hydrofluoric acid HBr (aq) (bromide ion) Hydrobromic acid HI (aq) (iodide ion) Hydroiodic acid HC2H3O2 (aq) (acetate ion) Acetic Acid HCN (aq) (cyanide ion) Hydrocyanic acid H2CO3 (aq) (carbonate ion) Carbonic Acid 3. Writing formulas: Oxidation numbers & polyatomic ions Formulas with Oxidation Numbers (pgs. 123, examples 124 ) To keep track of the transfer of electrons in all formulas, chemists have devised a system of electron bookkeeping called oxidation states (or oxidation numbers). In this method, an oxidation state is assigned to each member of a formula or polyatomic ion, It is designed by a small, whole-number superscript
preceded by a plus or minus sign. This is not to be confused with the ionic charges we have been using thus far that are used to the right of the ionic charge. These charges are directly related to the bonding that occurs in compounds. (Oxidations states are also used to track electron transfers.) The atom of an electron w/ a stronger pull for electrons is assigned the more negative charge. Rules for Assigning an Oxidation State like little sheet o Below are the basic rules for assigning an oxidation state to each element. By applying simple rules, oxidation states can be assigned to most elements or compounds. To apply these rules, remember that the sum of the oxidation states must be zero for an electrically neutral compound. For an ion, the sum of the oxidation states must equal the charge of the ion. The oxidation state of. 1. An atom in an element is zero. Examples: O for Na(s), O2(g), and H(l) 2. A monoatomic ion is the same as its charge. Examples: Na +1, Cl -1 3. Oxygen is usually -2 in its compounds. Example: H2O where 2H(+1) + 1O(-2) = 0. (Exceptions occur such as peroxide (O2 ) when the oxidation state/# is -1. 4. Hydrogen is usually +1. Examples: H2O, HCl, NH3 (The exception is in binary metal hydride compounds like NaH or CaH2.) 5. In binary compounds, the more electronegative element is assigned an oxidation number equal to its normal anion charge. Examples: PF3 : F=-1 and P2S3 ; S=-2 Formulas with Polyatomic Ions (page 121 for examples) When writing formulas using polyatomic ions, the rules do not change. Simply treat the polyatomic ion as if it were a single anion. If the cation is from category I, follow the rule for category I. If the cation is from category II, follow the rules for category II. The crisscross method does not change either. B. Percent Composition from Formulas 1. Formula masses- p.126 bottom By using the atomic masses assigned to the elements, we can find the formula mass of the compound.
Determined by multiplying atomic mass units (amu) as a whole number by the subscript for that element and then adding these values for all the elements in the formula. Example: Ca (OH)2 (one calcium amu + two hydrogen and two oxygen amu = formula mass). 1Ca (amu=40) = 40 2O (amu=16) = 32 2H (amu=1) = 2.0 Formula mass Ca (OH)2 = 74 amu If you have 6.02 x 10 23 atoms of an element, then the atomic mass units can be expressed in grams, and then the formula mass can be called the molar mass. (pgs. 164-165 explanation) o Another example if Fe2O3 2Fe (amu=56) = 112 3O (amu=16) = 48.0 Formula mass Fe2O3 = 160 amu 2. Calculations Percentage Composition Pages 127 and 128-129 for examples It is sometime useful to know what percent of the total weight of a compound is made up of a particular element. This is called finding the percentage composition. This simple formula for this is: Total amu of the element in the compound 100% = Percentage composition Total formula amu of that element To find the percent composition of calcium in calcium hydroxide, we set the formula up as follows: Ca = 40 amu 100% = 54% Calcium Formula mass = 74 amu To find the percent composition of oxygen in calcium hydroxide: O = 32 amu 100% = 43% Calcium Formula mass = 74 amu To find the percent composition of hydrogen in calcium hydroxide: H = 2.0 amu 100% = 2.7% Calcium Formula mass = 74 amu
Note - Hydrates Hydrates water (H2O) bound to ionic crystal / a compound that has water molecules included in its crystalline makeup When writing it as part of the symbol in chemical compounds you put a raised period before it (like a multiplication sign), and then whatever number the prefix is and then H2O next to that. When writing the name put the prefix+hydrate. Examples o CuSo 4 5H 2 O Copper (II) Sulfate Pentahydrate o NiSo 4 7H 2 O Nickel (II) Sulfate Heptahydrate