CHEMISTRY 1: Chemical Bonding Brown, T.L., LeMay, H.E., and B.E. Bursten. Chemistry: The Central Science 8 th Ed. New Jersey: PrenticeHall, 2002. CHEMICAL BONDS The forces that hold atoms or ions together in compounds are called chemical bonds. Chemical bonds can be classified as ionic, covalent, or metallic. Atoms tend to lose, gain, or share enough electrons to achieve a noblegas electron configuration. These observations are summarized in the octet rule. Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons. IONIC BONDING Many compounds consist of ions rather than molecules. Such compounds are said to be ionic. An ion is an electrically charged "package" consisting of one (monatomic ion) or more (polyatomic ion) atoms. An ion with a positive charge is called a cation (CATion), while an ion with a negative charge is called an anion (ANion). Ionic compounds do not exist as molecules and so do not have molecular formulas. Rather, ionic substances such as sodium chloride and magnesium chloride have only empirical formulas NaCl and MgCl 2, respectively. The charges on many atomic ions can be predicted using the periodic table. In general, for a nonmetal to form an ion, it will gain as many electrons as it needs in order to have the same number of electrons as a noble gas. Metals will lose electrons to become cations, while nonmetals will gain electrons to become anions. NAMING IONIC COMPOUNDS Listed in the following table are some of the common ions you need to know. The list seems long, but actually, you only need to memorize a few species. For other species in the main periodic table groups, you can deduce their charges by looking at the periodic table. COMMON CATIONS COMMON ANIONS CHARGE FORMULA NAME CHARGE FORMULA NAME H+ hydrogen ion or proton H hydride Li + lithium ion F fluoride Na + sodium ion Cl chloride 1+ K + potassium ion Br bromide Cs + cesium ion I iodide Ag + silver ion CN cyanide NH4+ ammonium ion 1 OH hydroxide Mg 2+ magnesium ion C2H3O2 acetate Ca 2+ calcium ion ClO3 chlorate Sr 2+ strontium ion ClO4 perchlorate Ba 2+ barium ion NO3 nitrate Zn 2+ zinc ion MnO4 permanganate Cd 2+ cadmium ion N3 azide Co 2+ cobalt (II) ion O 2 oxide 2+ Cu 2+ copper (II) O2 2 peroxide Fe 2+ iron (II) S 2 sulfide Mn 2+ manganese (II) 2 CO3 2 carbonate Hg22+ mercury (I) CrO4 2 chromate Hg 2+ mercury (II) Cr2O7 2 dichromate Ni 2+ nickel (II) SO4 2 sulfate Pb 2+ lead (II) 3 N 3 nitride Sn 2+ tin (II) PO4 3 phosphate Al 3+ aluminum ion 3+ Cr 3+ chromium (III) 1 HCO3 Fe 3+ iron (III) bicarbonate
SOME MORE RULES IN NAMING IONIC COMPOUNDS 1. Cations If the metal can form more than one cation, then the charge is indicated (in Roman numerals) in parentheses in the name. Examples: Cu + = copper(i) ion; Cu 2+ = copper(ii) ion Fe 2+ = iron(ii) ion; Fe 3+ = iron(iii) ion Older nomenclature system uses the ending ous to cation with fewer + charge and ic to the one with more + charge. Examples. Fe 2+ = ferrous ion Fe 3+ = ferric ion Cu + = cuprous ion Cu 2+ = cupric ion Cations formed from nonmetals end in ium. Examples: NH + 4 ammonium ion H 3 O + = hydronium ion 2. Anions Monoatomic anions have names formed by replacing the ending of the name of the element with ide. Examples: H = hydride ion N 3 = nitride ion A few simple polyatomic anions also have names ending in ide Examples: OH = hydroxide ion O 2 2 = peroxide ion CN = cyanide ion Polyatomic ions with oxygen have names ending in ate or ite Examples: ate used for most common oxyanion of an element ite used for oxyanion with same charge, but one less oxygen Examples: NO 3 = nitrate ion SO 2 4 = sulfate ion NO 2 = nitrite ion SO 2 3 = sulfite ion Prefixes are used when a series of oxyanions extends to 4 members, like with the halogens Per one more O atom than oxyanion ending in ate Hypo one less O atom than oxyanion ending in ite Examples: ClO 4 perchlorate ion ClO 3 chlorate ion ClO 2 chlorite ion ClO hypochlorite ion IO 4 periodate ion IO 3 iodate ion Anions derived by adding H+ to an oxyanion are named by adding as a prefix the word hydrogen or dihydrogen as appropriate 2 Examples: CO 3 carbonate ion HCO 3 hydrogen carbonate ion (or bicarbonate ion) 3 PO 4 phosphate ion H 2 PO 4 dihydrogen phosphate ion 3. Hydrates compounds with a specific number of water molecules attached Examples CuSO 4 5 H 2 O copper (II) sulfate pentahydrate BaCl 2 2 H 2 O barium chloride dihydrate Sr(NO 3 ) 2 4 H 2 O strontium nitrate tetrahydrate
NAMING MOLECULAR COMPOUNDS Naming binary molecular compounds is very much like naming ionic compounds, with one difference. With molecular compounds, we use Greek prefixes to denote how many of each atom are present. The procedures used for naming binary (twoelement) molecular compounds are similar to those used for naming ionic compounds: 1. The name of the element farthest to the left in the periodic table is usually written first. 2. If both elements are in the same group in the periodic table, the lower one is named first. 3. The name of the second element is given an ide ending. 4. Greek prefixes (Table 2.6) are used to indicate the number of atoms of each element. The prefix mono is never used with the first element. When the prefix ends in a or o and the name of the anion begins with a vowel (such as oxide), the a or o is often dropped. (ex. Carbon Monoxide, dinitrogen pentoxide) 5. Some exceptions compounds with H with common names Examples. B 2 H 6 diborane CH 4 methane SiH 4 silane NH 3 ammonia NAMING ACIDS and BASES 1. Acids substance that gives hydrogen ions (H + ) in water, i.e. in aqueous solution (aq). Anions with names ending in ide form acids with a hydro prefix and an ic ending Ex. HCl (aq) hydrogen chloride becomes hydrochloric acid when in water HF hydrogen fluoride becomes hydrofluoric acid HI hydroiodic acid HCN hydrocyanic acid Anions with names ending in ate form acids with an ic ending and those ending in ite form acids with an ous ending Ex. HNO 3 nitric acid (from NO 3 nitrate) H 2 SO 4 sulfuric acid 2 (from SO 4 sulfate) H 2 SO 3 sulfurous acid 2 (from SO 3 sulfite) Oxoacids acids with H, O and another element usually written with H first, then the central element, then O Ex. HNO 3 nitric acid H 2 CO 3 carbonic acid H 2 SO 4 sulfuric acid HClO 3 chloric acid
Sometimes, 2 or more acids have the same central atom but different number of O atoms. How to name: Addition of 1 O to ic acid : per ic acid Ex. HClO 3 chloric acid HClO 4 perchloric acid Removal of 1 O from ic acid : ous acid Ex. HNO 3 nitric acid HNO 2 nitrous acid Removal of 2 O atoms from ic acid: hypo ous acid Ex. HBrO 3 bromic acid HBrO hypobromous acid Example: HClO 4 perchloric acid ClO 4 perchlorate HClO 3 chloric acid ClO 3 chlorate HClO 2 chlorous acid ClO 2 chlorite HClO hypochlorous acid ClO hypochlorite 2. Bases substance that gives hydroxide ions (OH) in water; follow the naming for ionic compounds. Ex. NaOH sodium hydroxide KOH potassium hydroxide Ba(OH) 2 barium hydroxide DRAWING LEWIS STRUCTURES Drawing Lewis structures is an important first step toward understanding bonding in compounds. Use the following guidelines to practice drawing these structures. If you learn to follow this sequence step be step, you ll generate the correct Lewis structure. 1. Sum the valence electrons from all atoms. (Use the periodic table to help you do this.) For an anion, add an electron to the total for each negative charge. For a cation, subtract an electron for each positive charge. Don't worry about keeping track of which electrons come from which atoms. Only the total number is important. 2. Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond (a dash, representing two electrons). You usually have to choose a central atom for this. Chemical formulas are often written in the order in which the atoms are connected in the molecule or ion. When a central atom has a group of other atoms attached to it, the central atom is usually written first. The central atom is also usually the biggest, the relatively nearest to the center of the periodic table, or the least electronegative in the system. 3. Complete the octets of the atoms bonded to the central atom. (Remember, however, that hydrogen can have only two electrons.) Then count the number of electrons drawn so far. 4. Place any leftover electrons on the central atom, even if doing so results in more than an octet. 5. If there are not enough electrons to give the central atom an octet, try multiple bonds. Use one or more of the unshared pairs of electrons in the atoms bonded to the central atom to form double or triple bonds. 6. Do one final check if the number of electrons is correct, and if the octet rule is observed. If the octet rule is not observed, check if it belongs to any of the known exceptions. MOLECULAR GEOMETRY Lewis structures and the valenceshell electronpair repulsion (VSEPR) model. This model is based on the idea that the best arrangement of a given number of electron pairs is the one that minimizes the repulsions among them. We will apply the model first to polyatomic molecules of the general formula AB n (n can equal 2 through 6), where A is the central atom and is bonded to 2, 3, 4, 5, or 6 B atoms. Carbon tetrachloride is an example of an AB 4 molecule where A is carbon and B is chlorine.
To determine the molecular geometry (or shape) of a molecule or polyatomic ion, we follow a fourstep procedure. 1. Sketch the Lewis structure of the molecule or ion. 2. Count the total number of electron domains around the central atom, and arrange them in the way that minimizes the repulsion among them. 3. Describe the molecular geometry in terms of the angular arrangement of the bonded atoms. 4. A double or triple bond is counted as one electron domain when predicting.
CH1 PROBLEM SET 1 (due: Sept 3) Name: Yr. & Course: 1. When given the chemical formula, write the name. When given the name, write the formula. [60 pts] a. Potassium hydroxide p.cshco 3 b. Ammonium nitrate q.nan 3 c. Cadmium nitride r.ba 3 N 2 d. Barium perchlorate s.cdcn e. Aluminum sulfate t. FeBr3 f. Chromium (III) sulfi u. Hg 2 Cl 2 g. Calcium biphosphate v. NH 4 H 2 PO 4 tetrahydrate h. Sodium acetate w. Al 2 (SO 4 ) 3 12 H 2 O i. Magnesium perchlorate x. HgF 2 j. Nickel (II) oxide y. Mg(OH) 2 k. Nickel (II) peroxide z.zncro 4 l. Ferrous dichromate aa. PbCr 2 O 7 m. Periodic acid ab. HCl (aq) n. Diborane ac. HI (g) o. Sodium hydrogen ad. HClO carbonate 2. Draw the complete Lewis structure for the following. Please indicate the electron pair geometry AND the molecular geometry. Please box or encircle your final answers. Use the back of this sheet. [30pts] a. SO 2 b. SO 4 2 c. PO 4 3 d. O 3 e. CH 2 Cl F f. XeF 4 3. This is the structure of fructose. Indicate the molecular geometry around the labeled atoms. All lone pairs are indicated. (Treat each labeled atom as the central atom of every species.) [12pts.] 1. 2. 3. 4. 5. 6.