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Chemical Formulas and Chemical Compounds
Chapter 7
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Chemical Formulas Combinations of symbols are used to represent compounds of two or more elements. Also indicate the ratio of the number of atoms of each type of element in the compound. H2O – means that there are 2 hydrogen atoms for every oxygen atom. No subscript on O – means there is 1 Chemistry chapter 7
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Chemical Formulas Show either one molecule or one formula unit
Chemistry chapter 7
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Organic Compounds Written differently than other formulas
The shorthand shows how the atoms are joined, not just the number present. Example – CH3COOH, not C2H4O2 Chemistry chapter 7
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Ions Ion – charged atom or group of atoms Monatomic Ions – single atom
Polyatomic Ions – more than one atom Chemistry chapter 7
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Monatomic Ions Can be anions or cations
Transition elements can form more than one kind of ion See table 7-1 on page 205 You must memorize this table. Chemistry chapter 7
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Naming monatomic ions Cations Anions Element’s name
Roman numerals are used when there are multiple ions Anions Drop the element name ending Add -ide Chemistry chapter 7
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Binary compounds Contain two different elements
When we write chemical formula for a compound, the charges must add up to zero. Write the positive ion first. Chemistry chapter 7
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Example Write a formula for a compound of tin (II) and Iodine.
Tin (II) is 2+ Iodine is 1- We need two iodines to cancel out the charge on the tin (II). SnI2 Chemistry chapter 7
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Nomenclature Naming system Works for most compounds
Chemistry chapter 7
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Naming binary compounds
Write the name of the positive cation first. Add the name of the negative anion AlN – Aluminum nitride KCl – potassium chloride Chemistry chapter 7
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The stock system Elements with more than one possible charge
Cu2S – copper (I) sulfide CuS – copper (II) sulfide Note – in an older naming system the above could be written as cuprous sulfide and cupric sulfide Chemistry chapter 7
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Oxyanions Polyatomic ions that contain oxygen
When there are two or more oxyanions formed from the same two elements, the most common has the ending –ate The ion with one less oxygen than –ate ends in –ite The ion with one less oxygen than –ite adds the prefix hypo- The ion with one more oxygen than –ate adds the prefix per- Chemistry chapter 7
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Compounds with polyatomic ions
See table 7-2 on page 210 They are written like binary compounds. Except the ending isn’t changed to end in –ide CuSO4 – copper (II) sulfate Sn(SO4)2 – tin (IV) sulfate Chemistry chapter 7
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Discuss Practice problems 7-1, 7-2, and 7-3 on pages 207, 209, and 211
Chemistry chapter 7
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Polyatomic ions you must memorize
Ammonium Acetate Chlorate Chlorite Hydroxide Hypochlorite Nitrate Nitrite Perchlorate Permanganate Carbonate Peroxide Sulfate Sulfite Phosphate Chemistry chapter 7
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Naming binary molecular compounds
Two systems – one will be covered in section 7-2 Older system Prefixes used – see table 7-3 on page 212 CO – carbon monoxide CO2 – carbon dioxide SO2 – sulfur dioxide SO3 sulfur trioxide Chemistry chapter 7
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Rules List the less-electronegative element first. The second element
Only has a prefix if there is more than one. The second element Has a prefix Root of the element name -ide ending If the word begins with a vowel, drop the o or a at the end of the prefix (monoxide, not monooxide) Order: C, P, N, H, S, I, Br, Cl, O, F Chemistry chapter 7
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Examples PF5 N2O5 OF2 Phosphorus pentafluoride Dinitrogen pentoxide
Oxygen difluoride Chemistry chapter 7
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Acids Have a different naming rules.
Some common ones are listed in table 7-5 on page 214 You should know Hydrochloric acid (HCl) Sulfuric acid (H2SO4) Acetic acid (CH3COOH) (vinegar) Chemistry chapter 7
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Salts An ionic compound composed of a cation and the anion from an acid Sometimes the salt keeps one or more hydrogen atoms from the acid The prefix bi- or the word hydrogen is added to the anion name HCO3- Hydrogen carbonate ion or bicarbonate ion Chemistry chapter 7
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Discuss Sample problem 7-4 on page 213 Practice Chemistry chapter 7
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Discuss Chemistry chapter 7
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Oxidation numbers Also called oxidation states
Assigned to atoms in molecules Indicate the general distribution of electrons among the bonded atoms Sort of like ionic charge Chemistry chapter 7
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Pure elements Have oxidation numbers of zero Single atoms – Na
Molecules of a pure substance O2 P4 S8 Chemistry chapter 7
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Like charges on ions Shared electrons are assumed to belong to the more-electronegative atom The more electronegative element gets a number equal to the negative charge it would have as an anion. The less electronegative element gets a number equal to the positive charge it would have as a cation. Chemistry chapter 7
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Fluorine Oxidation number of -1 The most electronegative element
Chemistry chapter 7
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Oxygen Usually -2 In peroxides, -1 In compounds with halogens, +2 H2O2
OF2 Chemistry chapter 7
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Hydrogen +1 with more electronegative elements -1 with metals
Chemistry chapter 7
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Sum of oxidation numbers
In a neutral compound must be zero In a polyatomic ion must equal the charge on the ion Chemistry chapter 7
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Ion Can be assigned an oxidation number equal to the charge on the ion
Chemistry chapter 7
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Example Assign oxidation numbers to each atom in the following compound: KClO4 O is -2, which gives -8, since there are 4. The charge on perchlorate is 1-, so Cl must be +7 K must be +1 to cancel out the 1- +1, +7, -2 Chemistry chapter 7
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Example Assign oxidation numbers to each atom in the following compound: SO32- O is -2, which gives -6, since there are 3. The charge on sulfite is 2-, so S must be +4 +4, -2 Chemistry chapter 7
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You try Assign oxidation numbers to each atom in the following compound: CO2 O is -2, which gives -4, since there are 2. The charge is 0, so C must be +4 +4, -2 Chemistry chapter 7
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You try Assign oxidation numbers to each atom in the following compound: NO3- O is -2, which gives -6, since there are 3. The charge is 1-, so N must be +5 +5, -2 Chemistry chapter 7
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More oxidation numbers
See Appendix Table A-15 There is also a pattern on the periodic table Group 1 is usually +1 Group 2 is usually +2 Group 13 is usually +3 Group 14 is usually +2 or +4 Group 15 is usually -3 Group 16 is usually -2 Group 17 is usually -1 Chemistry chapter 7
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The stock system Can be used instead of prefixes for molecular compounds Use the oxidation number SO2 Sulfur dioxide Sulfur (IV) oxide SO3 Sulfur trioxide Sulfur (VI) oxide Chemistry chapter 7
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Discuss Name each of the following binary molecular compounds according to the stock system CI4 SO3 As2S3 NCl3 Chemistry chapter 7
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Formula mass The sum of the average atomic masses of all the atoms in a formula For ionic compounds or molecules Can also be called molecular mass for molecules Chemistry chapter 7
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Example Find the formula mass of Na2SO3 amu Chemistry chapter 7
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Example Find the formula mass of HClO3 84.46 amu Chemistry chapter 7
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You try Find the formula mass of MnO4- amu Chemistry chapter 7
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You try Find the formula mass of C2H6O 46.08 amu Chemistry chapter 7
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Molar Mass Chapter 3 The mass in grams of one mole (6.022 x 1023 particles) of a substance Example: H2O The mass of two moles of hydrogen atoms and one mole of oxygen atoms Chemistry chapter 7
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Example Find the molar mass of K2SO4 g/mol Chemistry chapter 7
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You try Find the molar mass of (NH4)2CrO4 152.10 g/mol
Chemistry chapter 7
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Formula mass and molar mass
Numerically equal Only the units are different Chemistry chapter 7
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Discuss How many moles of atoms of each element are there in one mole of ammonium carbonate, (NH4)2CO3 2 mol N, 8 mol H, 1 mol C, 3 mol O Determine both the formula mass and the molar mass of ammonium carbonate 96.11 amu, g/mol Chemistry chapter 7
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Converting with molar mass
Relate mass in grams to number of moles Relate mass in grams to number of particles Chemistry chapter 7
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Example What is the mass in grams of 3.04 mol of ammonia vapor, NH3?
Chemistry chapter 7
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You try What is the mass in grams of mol of calcium nitrate, Ca(NO3)2? 42.2 g Chemistry chapter 7
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Example How many moles of SO2 are in 3.82 g? 0.0596 mol
Chemistry chapter 7
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You try How many moles of Cl2 are there in 77.1 g? 1.09 mol
Chemistry chapter 7
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Example How many molecules are there in 77.1 g Cl2?
6.55 x 1023 molecules Chemistry chapter 7
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You try How many molecules are in 4.15 x 10-3 g of C6H12O6?
1.39 x 1019 molecules Chemistry chapter 7
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Percentage composition
Percentage by mass of each element in a compound Example: gum Chemistry chapter 7
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Example Find the percentage composition of sodium nitrate NaNO3.
Chemistry chapter 7
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You try Find the percentage composition of silver sulfate, Ag2SO4.
Chemistry chapter 7
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Discuss Zinc chloride, ZnCl2 is 52.02% chlorine by mass. What mass of chlorine is contained in 80.3 g of ZnCl2? How many moles of Cl is this? 41.8 g 1.18 mol Chemistry chapter 7
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Empirical formula The symbols for the elements combined in a compound
Subscripts show the smallest whole-number mole ratio of the atoms Determined from the percent composition of a substance Chemistry chapter 7
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Empirical formula Usually the same as an ionic compound’s formula unit
Not always the same as the molecular formula Diborane’s molecular formula is B2H6 The empirical formula is BH3 Chemistry chapter 7
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Example A compound is analyzed and found to contain 36.70% potassium, 33.27% chlorine, and 30.03% oxygen. What is the empirical formula of the compound? KClO2 Chemistry chapter 7
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You try Determine the empirical formula of the compound that contains 17.15% carbon, 1.44% hydrogen, and 81.41% fluorine. CHF3 Chemistry chapter 7
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Example A 60.0 g sample of tetraethylead, a gasoline additive, is found to contain g lead, g carbon, and 3.74 g hydrogen. Find its empirical formula PbC8H20 Chemistry chapter 7
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You try A g sample of an unidentified compound contains g sodium, g chromium, and g oxygen. What is its empirical formula? Na2Cr2O7 Chemistry chapter 7
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Discuss Find the empirical formula of a compound that contains 53.70% iron and 46.30% sulfur. Fe2S3 Chemistry chapter 7
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Molecular formula Show how many atoms are in each molecule
Related to empirical formula x is the whole number the subscripts must be multiplied by It might be 1 Chemistry chapter 7
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Mass relationship Chemistry chapter 7
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Example The empirical formula for trichloroisocyanuric acid is OCNCl. The molar mass of this compound is g/mol. What is its molecular formula? O3C3N3Cl3 Chemistry chapter 7
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Example Determine the molecular formula of a compound with an empirical formula of NH2 and a formula mass of amu. N2H4 Chemistry chapter 7
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You try Determine the molecular formula of the compound with an empirical formula of CH and a formula mass of amu. C6H6 Chemistry chapter 7
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Example If 4.04 g of N combine with g of O to produce a compound with a formula mass of amu, what is the molecular formula of this compound? N2O5 Chemistry chapter 7
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You try The molar mass of a compound is 92 g/mol. Analysis of a sample of the compound indicates that it contains g N and g O. Find its molecular formula. N2O4 Chemistry chapter 7
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