Chemistry Chapter 7 Chemical Names and FormulasPresentation Transcript
Chapter 7 Chemical Formulas and Chemical Compounds
Section 1 – Chemical Names and Formulas
Significance of Chemical Formula Chemical formula shows number and types of atoms in compound C8H18 Subscript indicates that there are 8 carbon atoms in a molecule of octane. Subscript indicates that there are 18 hydrogen atoms in a molecule of octane.
Ionic compound made of mixture of + and – ions held by attraction Combination of cation and anion Al2(SO4)3 Subscript 3 refers to everything inside parentheses giving 3 sulfate ions, with a total of 3 sulfur atoms and 12 oxygen atoms.
Monoatomic Ions Monoatomic ions – ions formed from single atom Group 1 – lose one e-, form +1 Group 2 – lose two e-, form +2 Groups 15, 16, 17 gain electrons to form anions -1, -2, -3
Not all main-group elements form ions easily C and Si form covalent bonds Transition metals can form +1, +2, +3, +4
Naming Monoatomic Ions Cations named same as element name K+ = potassium Mg2+ = magnesium Anions named by dropping ending and adding –ide F- = fluoride N3- = nitride
Binary Ionic Compounds Binary compounds compounds made of two different elements Total positive and negative charge must be equal Mg2+ Br- +2 + -1 = +1 +2 + 2(-1) = 0 Formula = MgBr2
Naming Binary Ionic Compounds Nomenclature naming system Combine names of cations and anions Cation name ALWAYS comes first Then anion name Al2O3 Aluminum oxide
Practice Problem 1 Write formulas for the binary ionic compounds formed between the following elements: a. potassium and iodine KI b. magnesium and chlorine MgCl2 c. sodium and sulfur Na2S
d. aluminum and sulfur Al2S3 e. aluminum and nitrogen AlN
Practice Problem 2 Name the binary ionic compounds: a. AgCl Silver chloride b. ZnO Zinc oxide c. CaBr2 Calcium bromide
d. SrF2 Strontium fluoride e. BaO Barium oxide f. CaCl2 Calcium chloride
Stock System Some elements form two or more cations with different charges Stock system uses Roman numerals to show ion’s charge Roman numeral included in () right after metal name Fe2+ Iron (II)
CuCl2 Copper (II) chloride Roman numeral indicating charge Name of anion Name of cation
Practice Problem 1 Write the formula and give the name for the compounds formed between the following ions: a. Cu2+ and Br− CuBr2, copper(II) bromide b. Fe2+ and O2− FeO, iron(II) oxide c. Pb2+ and Cl− PbCl2, lead(II) chloride
d. Hg2+ and S2− HgS, mercury(II) sulfide e. Sn2+ and F− SnF2, tin(II) fluoride f. Fe3+ and O2− Fe2O3, iron(III) oxide
Polyatomic Ions Polyatomic ion ion containing more than one atom Most are oxyanions polyatomic ions that contain oxygen NO2- NO3- Ion with more oxygen ends in –ate (nitrate) Ion with less oxygen ends in –ite (nitrite)
Hypo- and Hyper- One oxyanion family has 4 members ClO- ClO2- ClO3- ClO4- Hypochlorite Chlorite ChloratePerchlorate Naming same as binary ionic compounds Cation first, anion second
Practice Problem 1 Write formulas for the following ionic compounds: a. sodium iodide NaI b. calcium chloride CaCl2 c. potassium sulfide K2S d. lithium nitrate LiNO3
e. copper(II) sulfate CuSO4 f. sodium carbonate Na2CO3 g. calcium nitrite Ca(NO2)2 h. potassium perchlorate KClO4
Practice Problem 2 Give the names for the following compounds: a. Ag2O silver oxide b. Ca(OH)2 calcium hydroxide c. KClO3 potassium chlorate
d. NH4OH ammonium hydroxide e. FeCrO4 iron(II) chromate f. KClO potassium hypochlorite
Naming Binary Molecular Compounds Molecular compounds made of covalently bonded molecules Usually happens between 2 NON-metals
Prefix System Uses prefixes to show how many atoms present
Examples CCl4 carbon tetrachloride P4O10 tetraphosphorousdecoxide There are rules for determining which ion comes first
Rules The less-electronegative element is given first. If only 1 atom, no prefix used Second element named by combining (a) prefix, (b) element root name, (c) –ide The o or a at the end of prefix dropped when word following beings with vowel
Practice Problem Name the following binary molecular compounds: a. SO3 sulfur trioxide b. ICl3 iodine trichloride c. PBr5 phorphorouspentabromide
Acids Most acids either binary acids or oxyacids Binary acids acids that are made of 2 elements Usually H and one of the halogens Oxyacids acids that contain H, O, and a 3rd element (usually nonmetal)
Naming Acids Binary acids – hydro(element root)icacid Ex. HCl – hydrochloric acid Oxyacids – with less oxygen, (element root)ous acid – with more oxygen, (elementroot)ic acid Ex. HNO2 = nitrous acid HNO3 = nitric acid
Salts Salt an ionic compound made of a cation and an anion from an acid Table salt – NaCl – contains anion from HCl Some salts have anions where one or more H atoms from acid are kept Named by adding hydrogen OR prefix bi- to anion name HCO3- hydrogen carbonate ion bicarbonate ion
Practice Problem 1 Write formulas for the compounds formed between the following: a. aluminum and bromine AlBr3 b. sodium and oxygen Na2O c. magnesium and iodine MgI2 d. Pb2+ and O2− PbO
e. Sn2+ and I− SnI2 f. Fe3+ and S2− Fe2S3 g. Cu2+ and NO3− Cu(NO3)2 h. NH4 + and SO42− (NH4)2SO4
Practice Problem 2 Name the following compounds using the Stock system: a. NaI sodium iodide b. MgS magnesium sulfide c. CaO calcium oxide
d. K2S potassium sulfide e. CuBr copper (I) bromide f. FeCl2 iron (II) chloride
Practice Problem 3 Write formulas for each of the following compounds: a. barium sulfide BaS b. sodium hydroxide NaOH c. lead(II) nitrate Pb(NO3)2 d. potassium permanganate KMnO4
e. iron(II) sulfate FeSO4 f. diphosphorus trioxide P2O3 g. disulfur dichloride S2Cl2 h. carbon diselenide CSe2
i. acetic acid CH3COOH j. chloric acid HClO3 k. sulfurous acid H2SO3 l. phosphoric acid H3PO4
Section 2 – Oxidation Numbers The charges on the ions making an ionic compound reflect the electron distribution of the compound. In order to indicate the general distribution of electrons among the bonded atoms in a molecular compound or a polyatomic ion, oxidation numbers, also called oxidation states, are assigned to the atoms composing the compound or ion. Unlikeionic charges, oxidation numbers do not have an exact physical meaning. However, oxidation numbers are useful in naming compounds, in writing formulas, and in balancing chemical equations. And they are helpful in studying certain types of chemical reactions.
Assigning Oxidation Numbers Shared electrons assumed to belong to more-electronegative atom in each bond
Rules 1. Atoms in a pure element have Ox# of zero. Na, O2, P4 = 0 2. More-electronegative element in binary molecular compound assigned the number equal to negative charge it has as ion. Less electronegative assigned equal to positive charge it has as ion. 3. Fluorine always is -1 because it is most electronegative.
4. Oxygen has -2 in almost all compounds, except Peroxides (H2O2, O = -1) Compounds with halogens (O = +) 5. H = +1 in all compounds with more electronegative element, H = -1 in compounds with metals 6. Sum of Ox# of all atoms in neutral compound = zero 7. Sum of Ox# of all atoms in polyatomic ion = charge of ion 8. Ox# of ions in ionic compounds = its charge
Sample Problem 1 Assign Ox# to each atom in UF6. Start by placing known Ox#s above appropriate elements. From rules, we know F is always -1
Multiply known Ox#s by appropriate number of atoms and place totals underneath matching elements. There are 6 F atoms, 6 x -1 = -6
UF6 is molecular According to guidelines, sum of Ox#s must be zero Total positive Ox#s must be +6
Divide total positive Ox#s by number of atoms +6 ÷ 1 = +6
H2SO4 +1-2 H2 S O4 +2-8 +1 +6 -2 H2 S O4 +2 +6 -8
Practice Problem Assign oxidation numbers to each atom in the following compounds or ions: a. HCl +1, -1 b. CF4 +4, -1 c. PCl3 +3, -1
d. SO2 +4, -2 e. HNO3 +1, +5, -2 f. KH +1, -1 g. P4O10 +5, -2
h. HClO3 +1, +5, -2 i. N2O5 +5, -2 j. GeCl2 +2, -1
Stock System Remember there were two ways to name covalent compounds Stock system Prefix system Some nonmetals can have more than one oxidation number Listed in Table A-15 in handout
Can use Roman numerals to show oxidation number
Practice Problem 1 Assign oxidation numbers to each atom in the following compounds or ions: a. HF +1, -1 b. CI4 +4, -1 c. H2O +1, -2 d. PI3 +3, -1
e. CS2 +4, -2 f. Na2O2 +1, -1 g. H2CO3 +1, +4, -2 h. NO2− +3, -2
i. SO42− +6, -2 j. ClO2− +3, -2 k. IO3− +5, -2
Practice Problem 2 Name each of the following binary molecular compounds according to the Stock system: a. CI4 carbon(IV) iodide b. SO3 sulfur(VI) oxide c. As2S3 arsenic(III) sulfide d. NCl3 nitrogen(III) chloride
Section 3 – Using Chemical Formulas As you have seen, a chemical formula indicates the elements as well as the relative number of atoms or ions of each element present in a compound. Chemical formulas also allow chemists to calculate a number of characteristic values for a given compound. In this section, you will learn how to use chemical formulas to calculate the formula mass, the molar mass, and the percentage composition by mass of a compound.
Formula Masses Like atoms, a molecule, formula unit, or ion has an average mass (amu) Simply add amu of all atoms in molecule H2O Amu H = 2 x 1.01 amu Amu O = 1 x 16.00 amu Amu H2O = 18.02 amu
Mass of water molecule can be named molecular mass Mass of 1 NaCl is not molecular mass b/c it’s not a molecule, it’s an ionic compound Formula mass sum of the average atomic masses of all the atoms represented in its formula
Practice Problem Find the formula mass of each of the following: a. H2SO4 b. Ca(NO3)2 c. PO43− d. MgCl2
a. H2SO498.08 amu b. Ca(NO3)2 164.10 amu c. PO43− 94.97 amu d. MgCl2 95.21 amu
Molar Masses Molar mass is the same are formula mass, but units are in grams/mole (g/mol) Molar Mass H2O = ? H = 2 x 1.01 g/mol O = 1 x 16.00 g/mol H2O = 18.02 g/mol
Practice Problem Find the molar mass of each of the compounds a. Al2S3 b. NaNO3 c. Ba(OH)2
a. Al2S3 150.17 g/mol b. NaNO385.00 g/mol c. Ba(OH)2 171.35 g/mol
Molar Mass as Conversion Factor g/mol can be used as a conversion factor to change from moles to mass What is the mass in grams of 2.50 mol of oyxgen gas?
Practice Problem 1 Ibuprofen, C13H18O2, is the active ingredient in many nonprescription pain relievers. Its molar mass is 206.29 g/mol. a. If the tablets in a bottle contain a total of 33 g of ibuprofen, how many moles of ibuprofen are in the bottle? b. How many molecules of ibuprofen are in the bottle? c. What is the total mass in grams of carbon in 33 g of ibuprofen?
a. 0.16 mol C13H18O2 b. 9.6 x 1022 molecules C13H18O2 c. 25 g C
Practice Problem 2 How many moles of compound are there in the following? a. 6.60 g (NH4)2SO4 b. 4.5 kg Ca(OH)2 a. 0.0500 mol b. 61 mol
Practice Problem 3 How many molecules are there in the following? a. 25.0 g H2SO4 b. 125 g of sugar, C12H22O11 a. 1.53 × 1023 molecules b. 2.20 × 1023 molecules
Percent Composition Percentage composition percentage by mass of each element in a compound Divide mass of element in sample of compound by total mass of sample, multiply by 100
Mass percentage of element in compound same regardless of sample’s size Simpler way to calculate…. Determine how many grams of element are present in 1 mole of compound Then divide this value by molar mass of compound, multiply by 100
Sample Problem Find the percentage composition of copper(I) sulfide, Cu2S.
1. Analyze Given: formula, Cu2S Unknown: percentage composition of Cu2S
2. Plan formula -> molar mass -> mass percentage of each element The molar mass of the compound must be found Then the mass of each element present in one mole of the compound is used to calculate the mass percentage of each element.
3. Compute Molar mass: Cu = 2 x 63.55 g Cu = 127.1 g Cu S = 1 x 32.07 g S = 32.07 g S 159.2 g/mol Cu2S
Sample Problem 2 As some salts crystallize from a water solution, they bind water molecules in their crystal structure. Sodium carbonate forms such a hydrate, in which 10 water molecules are present for every formula unit of sodium carbonate. Find the mass percentage of water in sodium carbonate decahydrate, Na2CO3•10H2O, which has a molar mass of 286.14 g/mol.
1. Analyze Given: chemical formula, Na2CO3•10H2O molar mass of Na2CO3•10H2O Unknown: mass percentage of H2O
2. Plan chemical formula -> mass H2O per mole of Na2CO3•10H2O -> % water The mass of water per mole of sodium carbonate decahydrate must first be found. This value is then divided by the mass of one mole of Na2CO3•10H2O.
3. Compute 1 mol Na2CO3•10H2O contains 10 mol of H2O
4. Evaluate Checking shows that the arithmetic is correct and that units cancel as desired.
Section 4 – Determining Chemical Formulas
When a new substance is synthesized or is discovered, it is analyzed to show its percentage composition From this data, the empirical formula is then determined An empirical formula consists of the symbols for the elements combined in a compound, with subscripts showing the smallest whole-number mole ratio of the different atoms in the compound
For an ionic compound, the formula unit is usually the compound’s empirical formula For a molecular compound, however, the empirical formula does not necessarily indicate the actual numbers of atoms present in each molecule For example, the empirical formula of the gas diborane is BH3, but the molecular formula is B2H6 In this case, the number of atoms given by the molecular formula corresponds to the empirical ratio multiplied by two.
Calculation of Empirical Formulas Determine empirical formula from percent composition Change percent composition to grams Ex. 29.9% H …. 29.9 g H Use molar mass to change grams to mole Divide all moles by smallest mole amount Result is the subscript for the empirical formula
Example Percent composition of diborane is 78.1% B and 21.9% H In 100.0 g sample of diborane, there is 78.1 g B and 21.9 g H
Mass composition of each element converted to composition in moles by dividing by molar mass
7.22 mol B to 21.7 mol H Not a ratio of smallest whole numbers Divide each number of mol by smallest number in ratio Rounding, empirical formula is BH3
Sample Problem 1 Quantitative analysis shows that a compound contains 32.38% sodium, 22.65% sulfur, and 44.99% oxygen. Find the empirical formula of this compound.
1. Analyze Given: percentage composition: 32.38% Na, 22.65% S, and 44.99% O Unknown: empirical formula
2. Plan percentage composition mass moles smallest whole-number mole ratio of atoms
3. Compute Mass composition (mass of each element in 100.0 g sample): 32.38 g Na, 22.65 g S, 44.99 g O
Divide each mole by smallest mole in ratio Na2SO4
4. Evaluate Calculating the percentage composition of the compound based on the empirical formula determined in the problem reveals a percentage composition of 32.37% Na, 22.58% S, and 45.05% O These values equal 100%
Practice Problem 1 Analysis of a 10.150 g sample of a compound known to contain only phosphorus and oxygen indicates a phosphorus content of 4.433 g. What is the empirical formula of this compound? P2O5
Practice Problem 2 A compound is found to contain 63.52% iron and 36.48% sulfur. Find its empirical formula. FeS
Practice Problem 3 Find the empirical formula of a compound found to contain 26.56% potassium, 35.41% chromium, and the remainder oxygen. K2Cr2O7
Practice Problem 4 Analysis of 20.0 g of a compound containing only calcium and bromine indicates that 4.00 g of calcium are present. What is the empirical formula of the compound formed? CaBr2
Calculating Molecular Formulas Empirical formula contains smallest whole number ratio Molecular formula is actual formula for compound Empirical CH Molecular C2H4 (ethene), C3H6 (cyclopropane)
Relationship between empirical and molecular: x(empirical formula) = molecular formula x is whole-number multiple Must know molecular mass Ex. Experiment shows molar mass of diborane (BH3) is 27.67 g/mol Empirical formula mass = 13.84 g/mol
Practice Problem 1 The empirical formula of a compound of phosphorus and oxygen was found to be P2O5. Experimentation shows that the molar mass of this compound is 283.89 g/mol. What is the compound’s molecular formula? P4O10
Practice Problem 2 Determine the molecular formula of the compound with an empirical formula of CH and a formula mass of 78.110 g/mol. C6H6
Practice Problem 3 A sample of a compound with a formula mass of 34.00 g/mol is found to consist of 0.44 g H and 6.92 g O. Find its molecular formula. H2O2