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Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith SIMPLE INORGANIC FORMULAS AND NOMENCLATURE Binary Molecules I. A binary molecule is formed when two nonmetals or metalloids combine. Electrons are shared so the bonding involved is known as CD VQ LJ:::N T bonding. 2. Sometimes these compounds have generic or common names (water) and they also have systemic names (dihydrogen monoxide). The common names must be memorized. The systemic name is more complicated but it has the advantage that the formula of the compound can be deduced from the name. 3. Simple binary compounds consist of only a few atoms. Systemic naming of these compounds follow the rules: • The elements, except for H, are written in order of increasing group number. • Prefixes are used to designate the number of each element present in the molecule. 4. The prefixes are: rl0IJO 6 HE)<A DI 7 HEPTA 3 II( \ 8 otm "-.ION 1-1 2 .~,.. . 4 TET12.A 9 5 Pi.::l'-iTA 10 5. Mono is assumed. ***** ;Jt'l./l::7L. used in front of the first element. If there is only one atom, the mono is Name the following binary molecules: (a) C02 (b) N203 (c) P4010 (d) DECIj N20 cr·W_mAl DIN DIOX.I])~ IT/106EN T!2IO'IJDl.:;- It:"Tf(l::j PHoSPHoR.US iJE(.o)(.1 D i2" birJ rTQ.0680 SV\otJo '1..1j)t;" Ionic Compounds I. Ionic compounds are formed between metals and nonmetals. The electrostatic force of attraction between the positive ion (cation) formed by the metal and the negative ion (anion) formed by the nonmetal is what holds the compound together. OVER r 2. The cation is named first and written first when writing a formula while the anion is both named and written second. 3. Cations can have one or more charges. These charges are known as oxidation numbers, or valences. 4. The transition metals and representative elements in Groups 13, 14, 15, and 16 have multiple oxidation numbers. 5. Roman numerals, enclosed in parentheses, are used after the name of the cation to designate the oxidation state of the cation - only if the cation has more than one positive oxidation state from which to choose. 6. This system of nomenclature is known as the Stock system. 7. Here are some simple rules that should help in the determination of the oxidation numbers of cations from the formulas of their compounds. (a) The oxidation number of any element in its free state (uncombined with other elements) is zero. Fe in a bar of iron is zero. 02 and N2 in the atmosphere are zero. (b) The oxidation number of the alkali metals in a compound is always 1+. (c) The oxidation number of the alkaline earth metals in a compound is always 2+. (d) Fluorine in a compound is always assigned an oxidation number of 1-. t. (e) The oxidation number of oxygen is almost always 2- in a compound. Exceptions to this rule would be peroxides, oi- where the oxidation number of each oxygen is 1-, and superoxides, Of where the oxidation number of each oxygen is Yz-. (0 In covalent compounds (with nonmetals) hydrogen is assigned an oxidation number of 1+ (examples are HCI, H20, NH3, CH4). (g) In metallic halides the halogen (F, CI, Br, I, and At) always has an oxidation number equal to 1-. (h) Sulfide, selenide, telluride, and polonide are always 2- in binary salts. (i) Nitrides, phosphides, and arsenides are always 3- in binary salts. U) All other oxidation numbers are assigned so that the sum of the oxidation numbers of each element equals the net charge on the molecule or polyatomic ion. In neutral compounds, the sum of the positive and negative charges must equal zero. ***** Determine the oxidation number of the underlined element: + (b) NaCI04 7 t I~ +3 \ Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith 7. Free elements, no matter how complex the molecule, have an oxidation number (valence or charge) equal to zero. OJ. Fe. He ~~ < T~ /1 1 L-td. 8. The following are diatomic or polyatomic elements in nature which must be committed to memory. These elements exist as neutral molecules in nature: (a) i3r,J. (g) r, (b) J:'l (h) 03 (c) J\J.1 (i) f (d) CJa (j) Sl? (e) I--b (k) C &0 (f) 0) (I) C/o y Charges and the Periodic Table 1. The periodic table can be used to help determine charges on many ions. Cations come from metals that lose electrons ( 0 -y: I IJ tCl, TI 0 J-J ) in order to become .ISO 8-ECTl20 tJ' C with a noble gas. Anions come from nonmetals that gain electrons ( lZ..Ebucno Iv ) in order to become r50E1£ CIlLo,.) f c.. with a noble gas. ***** Group 1 +1 Group 17 -I Group 2 +~ Group 16 - .) Group 3 +-3 Group IS -3 Group 14 -~ 2. Transition metals, representative metals with typically have more than one oxidation state. + and ~ s« +t..J sublevels, and the inner transition metals ei 3. Electrons for these metallic elements are lost in the following order: S) J.... . Such elements are ;...fa T isoelectronic with a noble gas when the outermost (valence) electrons are lost. 4. Inner transition elements are also known as the lanthanides, the actinides, the rare earth elements, and the transuranium elements. These elements are rare and many exist for short periods of time. 5. Both inner transition elements and transition elements are known for their variable oxidation numbers. The most common oxidation number for transition elements is f ~ OVER r 6. The d: sublevel in transition elements is responsible for the various oxidation numbers that result. Incomplete d sublevels are also responsible for the many colorful transition compounds that are known to exist. Complete d sublevels in cations of silver and zinc result in white compounds. SUMMARY OF CATIONS WITH VARIABLE OXIDATION NUMBERS - STOCK SYSTEM copper (I), Cu"; copper (II), Cu2+ mercury (I), Hgr"; mercury (II), Hg2+ 1+,2+ gold (I), Au"; gold (Ill), Au3+ indium 0), In"; indium (III), £n3+ thallium (0, TI+; thallium ([II), T13+ chromium (II), Cr2+;chromium (III), cr3+ cobalt (Il), C02+; cobalt (III), C03+ iron (II), Fez+; iron (III), Fe3+ lead (II), Pb2+; lead (IV), Pb4+ platinum (II), peT; platinum (IV), pt+ tin (II), Sn2+;tin (IV), Sn4+ zirconium (II), zirconium(lV), Zr4+ 1+,3+ 2+,3+ 2+,4+ z>. cerium (III), Ce3+; cerium (IV), Ce4+ 3+,4+ antimony (1II),Sb3\ antimony (V), Sb5+ arsenic (III), As3+; arsenic (V), As5+ bismuth(III), Bi3+; bismuth (V), Bi5+ phosphorus (III), p3+;phosphorus (V), p5+ iridium (II), Ir2+;iridium (III), Ir3+;iridium (IV), Ir4+ titanium (II), Ti2+;titanium (HI), Ti3+; titanium (IV), Ti4+ manganese (II), Mnl+; manganese (III), Mn3+;manganese (IV), Mn4+ 3+,5+ 2+,3+,4+ 2+,4+,5+ tungsten (II), W2+; tungsten (IV), W4+; tungsten (V), W5+ 3+,4+,5+ uranium (III), U3+;uranium (IV), U4+; uranium (V), U5+ 2+,3+,4+, 5+ vanadium (II), VZ+; vanadium (III), V3+;vanadium (IV), V4:; vanadium (V), V5+ - ! ' Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith Polyatomic Ions I. Polyatomic ions are a group of atoms that behave as a single ion. 2. The bonding within a polyatomic ion is covalent, but because there is either an excess or a shortage of electrons compared to the number of protons present, an ion results. 3. This short list of polyatomic ions must be MEMORIZED. NH4+ ammonium OR hydroxide NOf nitrite P043- phosphate N03- nitrate CIO- hypochlorite so> sulfite CIOf chlorite S042- sulfate C103• chlorate C032• carbonate CI04- perchlorate HC03- bicarbonate Cr042- chromate C2H)Of acetate Cr20l· dichromate CN- cyanide SCN- thiocyanate You also need to know the common Group ions such as (but not limited to) the ions in Groups 1 & 2, the halogens, 0, S, N, P, and "the triangle": Ag, Zn, Cd, AI, Ga, and In. ,. r \ OVER © Adrian Dingle's Chemistry Pages 2004. All rights reserved. These materials may NOT be copied or redistributed in any way, except for individual class instruction. Revised June 2005 AP Common Ions CAT10NS (+ve) Name Symbol ANIONS (-ve) Alternative" Name Symbol Alternatlve~ Bf Aluminum At'+ Bromide Ammonium NH: As3+ Bromate (I) BrO' (Hypobromite) Bromate (III) BrOz' As~ Bromate (V) BrO,' (Bromite) (Bromate) Barium Ba2+ Bromate (V") BrO•. (Perbromate) Bismuth (III) 8i3+ Carbonate C0,2' Arsenic (III) Arsenic (V) Bismuth (V) Bi~ Chlorate (I) Cadmium Calcium c<f+ Ca2+ " Chlomte (III) CIO' CI02' (Hypochlortte) (Chlorite) : Chlorate (V) CIO; (Chlorate) Chromium (II) Cr· ,Chlorate (VII) CIO" (Perchlorate) Chromium (III) Cr+ . Chloride Cobalt (II) C02+ Cobalt (III) Co'· Cu+ Chromate , Cyanide Copper (I) Hydrogen Cu2+ H+ Hydronium H,O+ Iron (II) Fe2+ Iron (III) Fe'+ ptf+ Copper (II) Lead (II) Lead (IV) cr crO.2, CN" CrzOr2, (Cuprous) Dichromate (Cupric) Dihydrogen Phosphate H2POi Ethanoate CzH,Oi F" Ruoride (Ferrous) (Ferric) Hydride (Plumbous) Hydrogen Oxalate Hydrogen Carbonate H' He03' (Acetate) (Bicarbonate) /: HCzO" HPO/' (Binoxalate) ( ,,Hydrogen Sulfate HSOi (Bisulfate ) ..Hydrogen Sulfide HS' HSO)' (Bisulfide) (Bisulfite) Hydrogen Phosphate \' Pb" L( Mg1+ Mn~+ Mn.•.• (Plumbic) Mercury (I) H!hz+ (Mercurous) Iodate (I) 10' (Hypoiodile) Mercury (II) Hg2• (Mercuric) Iodate (III) 102' (Iodite) Nickel (II) Ni2+ (Iodate) K'" Iodate (V) Iodate (VII) 10; Potassium 10. (Periodate) lithium Magnesium Manganese (II) Manganese (IV) Silver Na+ Strontium Sr+ Snz+ Tin (IV) Zinc Hydroxide Ag+ Sodium ,Tin (II) Hydrogen Sulrrte Sn'" Znz+ Iodide Manganate (V") OH' r MnO •. , Nitrote NO,' (Stannous) Nitride (Stannic) NiDite N"" NOz' Oxalate Oxide Peroxide Phosphate '~ CzO/' (Permanganate) (Ethandioate) ry- 0/' PO.3- Phosphide p"" Phosphite PO/' Sulfate SO.z, Sulfide Sulfite S1' 5032- Thiosulfate S20)2- Thiocyanate SCN" • In the case of the cations, the alternative names are genemlly redundant in modem chemistry, but the anions sometimes use the older, alternate names, For example, the oxyhatogen ions (bromate, chlorate, iodate stc.) are usually referred to by the altemate names, but HSOJ' is much more likely 10 be called Hydrogen Sulfrte rather than Bisufite, ,---,------- ~----- Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith Ternary Nomenclature: Acids and salts Containing Halogens and/or Oxygen 1. The halogens, with their variable oxidation numbers, allow for a great variety of compounds. !3/1~£" 2. A good way to learn ternary nomenclature is to start with a certain HOI'\<lE polyatomic ion. This is polyatomic ion ending with the suffix - C~ te Number of Oxygen Atoms (compared to home base) Polyatomic Ion Name Plus one oxygen atom Cl04' perchlorate ion Home base CI03' ch lorate ion Minus one oxygen atom CI02' chlorite ion Minus two oxygen atoms ClO' hypochlorite ion No oxygen atoms CI' chloride ion 3, Water solutions of binary hydrides form acids. The root derived from the hydride is given the prefix li'/ Dr2..0 and the suffix - Ie and the name ends with the word ACt D 4. The binary hydride HCl is known as H'I Dt2.0GeJ CI-ILo.e.IOlE' when aqueous it is known as 14 'f D1200-l W.e.IL !~ C ID (hydrogen monochloride) gas, but 5. Many common acids contain only oxygen, hydrogen, and a nonmetallic ion or polyatomic ion. These acids are called _-=O...:..Y--...:'1..:.,i-\,!.,:C:..;I.::::.b..=S=-_ 6, If the name of the polyatomic ion ends in the word acid. - ['L 7. If the name of the polyatomic ion ends in - / the word acid. 8, The mnemonic aid is: ***** (cfe- ic. +e , the suffix te. , the suffix ('fe - Name the following compounds: (a) HI04 (aq) (b) NaBr04 Sdb \v yt, PERIOble. I~( 11) eEQ8~Oi!~\(~TE OVER OU.5 1/ - f c..- is substituted followed by - OUSis substituted followed by ***** Write the formulas for the following compounds: (a) calcium hypochlorite Cl\. ( OD)::I (c) cyanic acid Ii (NO (b) hydrotelluric acid H.)Ie O!L iioe iJ (d) chlorous acid HC~O,;! BALANCING EQUATIONS HINTS SYNTHESIS & DECOMPOSITION Chemists write balanced equations to illustrate what is happening during a chemical reaction. Bonds are broken, atoms are rearranged, and new bonds are formed. Every chemical reaction supports the Law of conservation of Matter. This means that in every reaction, the number of atoms of each type of element contained within the reactants must be the same as the number of atoms of each type of element contained within the products. Balancing equations is a process which assures that equations are written properly to support the Law of Conservation of Matter; however, balancing cannot be done until each reactant and product formula is written correctly. It is important to properly write the seven elements that are diatomic in their elemental form and also use subscripts and parentheses appropriately when considering the oxidation number of ions. All compounds must be made neutral before beginning to balance the atoms. Balancing is accomplished by adding coefficients that multiply the number of atoms represented by the formula. For example, a coefficient of2 in front of oxygen (e.g., 2 02) means that 4 oxygen atoms are represented. Unlike algebra, in chemistry a coefficient does not need to be outside parentheses or brackets to b distributed. A coefficient applies to the complete substance; however, it no longer applies when a plus sign (+) or arrow (--+) is encountered. For example: 3 (NH4)2C03 shows 6 nitrogen, 24 hydrogen, 3 carbon, and 9 oxygen atoms. 3 MgCb + NaBr indicates 3 magnesium, 6 chlorine, 1 sodium, and 1 bromine atom. PREREQUISITE KNOWLEDGE Before you do anything, you must know and understand the following areas of nomenclature and formula writing: I. Ionic compounds 2. Covalent compounds 3. Acids and bases 4. Complex ions (coordination chemistry) 5. Organic nomenclature Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith TIPS FOR BALANCING EQUATIONS I. Ensure each molecular formula is written correctly and each compound is neutral. 2. Mentally count or tally how many of each type of atom is present on each side of the equation. 3. Begin by balancing elements that are only found in one substance on each side. 4. Balance oxygen and hydrogen LAST - they usually balance out at the end or perhaps only the number of water molecules needs to be adjusted. 5. If there is an odd number of an element on one side and an even number on the other, the odd number will need to be evened out - so use a coefficient of 2 for that substance. 6. Ifthere are polyatomic ions that remain together as a unit during the reaction, count the polyatomic ion as a unit. 7. When tallying, be sure to adjust the count for each and every element that an added coefficient affects. 8. Combustion reactions that don't seem to balance will often come out better if a coefficient of 2 is used for the hydrocarbon. SYNTHESIS REACTIONS Synthesis reactions occur when two or more reactants combine to form a single product. There are several types of synthesis reactions. I. A metal combines with a nonmetal to form a binary salt. Example: A piece of lithium metal is dropped into a container of nitrogen gas. 2. Metallic oxides and water form bases (metallic hydroxides) Example: Solid magnesium oxide is added to water. MgO + 2 HOH -+ Mg(OH)2 3. Nonmetallic oxides and water form acids. The nonmetal retains its oxidation number. Example: Dinitrogen pentoxide is bubbled into water. 4. Metallic oxides and nonmetallic oxides form salts. Example: solid calcium oxide is added to sulfur trioxide. CaO + S03 ~ CaS04 OVER DECOMPOSITION REACTIONS Decomposition reactions occur when a single reactant is broken down into two or more products. Ie > 1. Metallic carbonates decompose into metallic oxides and carbon dioxide. Example: A sample of magnesium carbonate is heated. 2. Metallic chlorates decompose into metallic chlorides and oxygen. Example: A sample of magnesium chlorate is heated. 3. Ammonium carbonate decomposes into ammonia, water and carbon dioxide. Example: A sample of ammonium carbonate is heated. 4. Sulfurous acid decomposes into sulfur dioxide and water. Example: A sample of sulfurous acid is heated. 5. Carbonic acid decomposes into carbon dioxide and water. Example: A sample of carbonic acid is heated. 6. A binary compound may break down into two elements. Example: Molten sodium chloride is electrolyzed. 2 NaCI 7. -+ 2 Na + Ch Hydrogen peroxide decomposes into water and oxygen. 8. Ammonium hydroxide decomposes into ammonia and water. " from The Ultimate Chemical Equations Handbook by Hague and Smith Re-typed NOMENCLA TURE PRACTICE I. Name each of the following compounds. (a) CaF2 (h) PFs Pt:NTt\ FU) ,:)H O:Sr l-\0LU.s F LUOe.1 Dt= CA Lt lIJ M ce. I DE (i) (NH4hS03 Si..'LFlll: 1~i\,nION'viH U) (c) NaH WI Dfll j) l: :30 b I Vi'Vl (m) GeLD (d) HIO (aq) H'IpcIO AU2(C204)3 (k) KSCN DoLlS {JOTAS S \ v r11 I~LI.t) (e) Hg2Ch '/l-11 OC'f A A.1 A n= (I) Pb(C2H302)4 n18Z.. C 1'Q.'( (I) C H L02. I O<i- LE'AD (t) Ba02 (IV) ACETATE (m) LiH i=t-e.o Y... I .D~ rM)'fLl U WI U T\-l h..H1I\ (g) Mn(HC03)3 rY\ i\NGArJESE (n) H2Cr04 (aq) ( ill.." H'I DI(OGe-J 2. Write formulas for the following (a) vanadium (V) oxide CAf2&iJ.RTE substances. (g) tin (IV) chromate 'Vd. O.s Sn~ (Cr OLl)'f (b) zinc hydroxide (h) nitrous acid l'1,(oH)J /f,JO(J (c) silver chromate Aj ,:J H '11)/2..\ DE (i) magnesium c-o; (d) tungsten (V) thiosulfate U) calcium carbide ~() (k) mercury (II) acetate C e (A103)-3 (t) hydroiodic acid phosphate Il1!J /-I1tJ C; ~J~(S.;l 03)S (e) cerium (III) nitrate hydrogen J/j (C,; 1-/3 Q;J) (I) perbromic (-1-1 acid !--I erO 'I OVER J BALANCING SYNTHESIS AND DECOMPOSITION EQUATIONS Predict and balance the following synthesis and decomposition reactions. Use abbreviations to indicate the phase of reactants and products where possible [(aq) (s) (I) (g)] I. A sample of calcium carbonate is heated. 2. Sulfur dioxide gas is bubbled through water. 3. Solid potassium oxide is added to a container ofcarbon dioxide gas. 4. Liquid hydrogen peroxide is warmed. 5. Solid lithium oxide is added to water. " L ., D "\ (5) /ti 0 to -'7 6. Molten aluminum chloride is electrolyzed. ----- ------------- Re-typed from The Ultimate Chemical Equations Handbook by Hague and Smith 7. A pea-sized piece of sodium is added to a container of iodine vapor . .;J J4.(s)+- I,J- Cj) -7 ,J Aft:-...-l. (5) 8. A sample of carbonic acid is heated. 9. A sample of potassium chlorate is heated. ') ..v L.Ji) AQ 1\ L..,( .3 (S) -~ d. k'tf (5) -I- .J 002 ~J) .. 10. Solid magnesium oxide is added to sulfur trioxide gas. OVER --~---- .. ---