1 Chapter 7 Chemical Formulas and Chemical Compounds
2 i) Understanding the Significance of a Chemical Formula ii) Monatomic Ions Identification and Nomenclature iii) Binary Ionic Compounds iv) Writing the Formula of an Ionic Compound v) Naming Binary Ionic Compounds vi) Naming Binary Molecular Compounds vii)covalent-network Compounds viii)acids and Salts
3 Objectives Chapter 7, Section 7.1 Chemical Names and Formulas i) Explain the significance of a chemical formula. ii) Determine the formula of an ionic compound formed between two given ions. iii) Name an ionic compound given its formula. iv) Using prefixes, name a binary molecular compound from its formula. v) Write the formula of a binary molecular compound given its name.
4 Significance of a Chemical Formula A chemical formula indicates the relative number of atoms of each kind in a chemical compound. For a molecular compound, the chemical formula reveals the number of atoms of each element contained in a single molecule of the compound. Example: Octane C 8 H 18 The subscript after the C indicates that there are 8 carbon atoms in the molecule. The subscript after the H indicates that there are 18 hydrogen atoms in the molecule.
5 Significance of a Chemical Formula The chemical formula for an ionic compound represents one formula unit = The simplest ratio of the compound s positive ions (cations) and its negative ions (anions). Example: aluminum sulfate Al 2 (SO 4 ) 3 Parentheses surround the polyatomic ion (SO 4-2 ) to identify it as a unit. The subscript 3 refers to the unit. Note also that there is no subscript for sulfur: when there is no subscript next to an atom, the subscript is understood to be 1.
6 Reading Chemical Formulas Octane C 8 H 18 Al 2 (SO 4 ) 3 Elemental Identity (Carbon) Elemental Identity (Hydrogen) Sulfate Polyatomic Ion w/ Subscript Indicating 3 Sulfates Elemental Identity (Aluminum)
7 Monoatomic Ions Many main-group elements can lose or gain electrons to form ions. Ions formed form a single atom are known as monatomic ions. Example: To gain a noble-gas electron configuration, nitrogen gains three electrons to form N 3- ions. Some main-group elements tend to form covalent bonds instead of forming ions. Examples are carbon and silicon
8 Naming Monoatomic Ions Monatomic cations are identified simply by the element s name. Examples: i) Element K 1e - K + is called the potassium cation ii) Element Mg 2e - Mg 2+ is called the magnesium cation For monatomic anions, the ending of the element s name is dropped, and the ending -ide is added to the root name. Examples: i) Element F + 1e - F - is called the fluoride anion ii) Element N + 3e - N -3 is called the nitride anion
10 Binary Ionic Compounds Compounds composed of two elements = binary compounds. Binary ionic compound the total numbers of positive charges and negative charges must be equal. The formula for a binary ionic compound can be written given the identities of the compound s ions. Example: Magnesium Bromide Chemical formula: MgBr 2 Ions combined: 1 Mg 2+ and 2 Br Mg 2+ + Br + Br MgBr 2
11 A general rule to use when determining the formula for a binary ionic compound is crossing over to balance charges between ions. Example: Aluminum oxide Chemical formula: Al 2 O 3 1) Write the symbols for the ions involved: Al 3+ and O 2-2) Cross over the charges by using the absolute value of each ion s charge as the subscript for the other ion. Al 3+ O 2 Al 2 O 3 Al 2 O 3 3) Check the combined positive and negative charges to see if they are equal. (2 3 + ) + (3 2 - ) = 0 Correct as Al 2 O 3
12 Writing the Formula of an Ionic Compound
13 Naming Binary Ionic Compounds The nomenclature, or naming system, of binary ionic compounds involves combining the names of the compound s positive and negative ions. i) Name of the cation is given first ii) Followed by the name of the anion Example: Al 2 O 3 > Aluminum oxide For most simple ionic compounds, the ratio of the ions is not given in the compound s name, because it is understood based on the relative charges of the compound s ions.
14 Sample Problem: Write the formulas for the binary ionic compound formed between the following elements: Zinc and Iodine Solution: 1) Write the symbols for the ions side by side. Write the cation first. Zn 2+ and I 2) Cross over the charges to give subscripts: Zn 2+ 1 I 2 3) Check the subscripts and divide them by their largest common factor to give the smallest possible whole-number ratio of ions. The subscripts give equal total charges of 1 2+ = 2+ and 2 1 = 2. The largest common factor of the subscripts is 1. The smallest possible whole-number ratio of ions in the compound is 1:2. The formula is ZnI 2
15 Sample Problem: Write the formulas for the binary ionic compound formed between the following elements: Zinc and Sulfur Solution: 1) Write the symbols for the ions side by side. Write the cation first. Zn 2+ and S 2-2) Cross over the charges to give subscripts: Zn S 2 3) The subscripts give equal total charges of 2 2+ = 4+ and 2 2 = 4. The largest common factor of the subscripts is 2. The smallest whole-number ratio of ions in the compound is 1:1. The formula is ZnS
16 Naming Binary Ionic Compounds The Stock System of Nomenclature Some elements such as iron, form two or more cations with different charges. To distinguish the ions formed by such elements, scientists use the Stock system of nomenclature. The system uses a Roman numeral to indicate an ion s charge. Examples: Fe 2+ iron(ii) Fe 3+ iron(iii)
17 The Stock System of Nomenclature Element Cr 2e - Cr +2 Element Cr 3e - Cr +3 Element Pb 2e - Pb +2 Element Pb 2e - Pb +4 Element Co 2e - Co +2 is called the Chromium (II) cation is called the Chromium (III) cation is called the Lead (II) cation is called the Lead (IV) cation is called the Cobalt (II) cation Element Co +2 2e - Co +4 is called the Cobalt (IV) cation
24 Naming Binary Ionic Compounds The Stock System of Nomenclature Example - Write the formula and give the name for the compound formed by the ions Cr 3+ and F -. Write the symbols for the ions side by side. Write the cation first. Cr 3+ F - Cross over the charges to give subscripts Cr 1 3+ F 3 The subscripts give charges of 1 3+ = 3+ and 3 1 = 3. The largest common factor of the subscripts is 1, so the smallest whole number ratio of the ions is 1:3 The formula is CrF 3 Chromium(III) Fluoride
25 Polyatomic Ions Many common polyatomic ions are oxyanions polyatomic ions that contain oxygen. Some elements can combine with oxygen to form more than one type of oxyanion. Examples: Nitrate NO 3 - and Nitrite NO 2 - The name of the ion with the greater number of oxygen atoms ends in -ate. The name of the ion with the smaller number of oxygen atoms ends in ite.
26 Polyatomic Ions Some elements can form more than two types of oxyanions. Example: chlorine can form ClO -, ClO 2-, ClO 3- or ClO 4 - In this case, an anion that has one fewer oxygen atom than the -ite anion has is given the prefix hypo-. An anion that has one more oxygen atom than the -ate anion has is given the prefix per-. Hypochlorite Chlorite Chlorate Perchlorate ClO - ClO - 2 ClO - 3 ClO - 4
27 Polyatomic Ions
28 Naming Compounds with Polyatomic Ions
29 Understanding Formulas for Polyatomic Ionic Compounds
30 Naming Compounds Containing Polyatomic Ions
31 Naming Compounds Containing Polyatomic Ions
32 Naming Compounds Containing Polyatomic Ions
33 Naming Compounds Containing Polyatomic Ions
34 Sample Problem: Write the formula for tin(iv) sulfate Solution: 1) Write the symbols for the ions side by side. Write the cation first. Sn 4+ and SO 4 2-2) Cross over the charges to give subscripts: Sn 2 4+ (SO 4 ) 4 2 3) The total positive charge is 2 4+ = 8+ The total negative charge is 4 2 = 8 4) The largest common factor of the subscripts is 2, so the smallest whole-number ratio of ions in the compound is 1:2 5) The correct formula is therefore Sn(SO 4 ) 2
35 Naming Binary Molecular Compounds Unlike ionic compounds, molecular compounds are composed of individual covalently bonded units, or molecules. As with ionic compounds, there is also a Stock system for naming molecular compounds. The old system of naming molecular compounds is based on the use of prefixes. Examples: i) CCl 4 Carbon tetrachloride (tetra = 4) ii) CO Carbon monoxide (mon = 1) iii) CO 2 Carbon dioxide (di = 2)
36 Prefixes for Naming Covalent Compounds
37 Prefixes for Naming Covalent Compounds
38 Naming Covalent Compounds
39 Naming Covalent Compounds
40 Naming Covalent Compounds
41 Naming Binary Molecular Compounds Sample Problem Give the name for As 2 O 5 The compound contains two arsenic atoms, so the first word in the name is diarsenic. The five oxygen atoms are indicated by adding the prefix pent- to the word oxide. The complete name is Diarsenic pentoxide.
42 Naming Binary Molecular Compounds Sample Problem Write the formula for oxygen difluoride. Oxygen is first in the name because it is less electronegative than fluorine. Because there is no prefix, there must be only one oxygen atom. The prefix di- in difluoride shows that there are two fluorine atoms in the molecule. The formula is OF 2
43 Covalent-Network Compounds Some covalent compounds do not consist of individual molecules. Instead, each atom is joined to all its neighbors in a covalently bonded, three-dimensional network. Subscripts in a formula for covalent-network compound indicate smallest whole-number ratios of the atoms in the compound. Examples: SiC Silicon Carbide SiO 2 Silicon Dioxide Si 3 N 4 Trisilicon Tetranitride.
44 Acids and Salts An acid is a certain type of molecular compound. Most acids used in the laboratory are either binary acids or oxyacids. Binary acids are acids that consist of two elements, usually hydrogen and a halogen. Oxyacids are acids that contain hydrogen, oxygen, and a third element (usually a nonmetal).
45 Acids and Salts In the laboratory, the term acid usually refers to a solution in water of an acid compound rather than the acid itself. example: Hydrochloric acid refers to a water solution of the molecular compound hydrogen chloride, HCl Many polyatomic ions are produced by the loss of hydrogen ions from oxyacids. Examples: Sulfuric Acid H 2 SO 4 Sulfate SO 2-4 Nitric Acid HNO 3 Nitrate NO - 3 Phosphoric Acid H 3 PO 4 Phosphate PO 3-4
46 Acids and Salts An ionic compound composed of a cation and the anion from an acid is often referred to as a salt. Examples: i) Table salt, NaCl, contains the anion from hydrochloric acid, HCl. ii) Calcium sulfate, CaSO 4, is a salt containing the anion from sulfuric acid, H 2 SO 4 iii) The bicarbonate ion, HCO 3, comes from carbonic acid, H 2 CO 3
47 Naming Binary Acids
48 Naming Oxyacids H 3 PO 4 H 2 SO 4
49 Naming Oxyacids HNO 3 H 2 CO 3
51 Salt Water Particle Diagram
52 Prefixes and Suffixes for Oxyanions and Related Acids
53 Objectives Chapter 7, Section 7.2 Oxidation Numbers i) List the rules for assigning oxidation numbers. ii) Give the oxidation number for each element in the formula of a chemical compound. iii) Name binary molecular compounds using oxidation numbers and the Stock system.
54 Oxidation Numbers Numbers called oxidation numbers can be assigned to atoms in order to keep track of electron distributions in molecular as well as ionic compounds. The charges on the ions in 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 are assigned to the atoms composing the compound or ion. Unlike ionic charges, oxidation numbers do not have an exact physical meaning: rather, they serve as useful bookkeeping devices to help keep track of electrons.
55 Assigning Oxidation Numbers In general when assigning oxidation numbers, shared electrons are assumed to belong to the more electronegative atom in each bond. More-specific rules are provided by the following guidelines. 1) The atoms in a pure element have an oxidation number of zero. Examples: all atoms in sodium, Na, oxygen, O 2, phosphorus, P 4, and sulfur, S 8, have oxidation numbers of zero. 2) The more-electronegative element in a binary compound is assigned a negative number equal to the charge it would have as an anion. Likewise for the less-electronegative element.
56 Assigning Oxidation Numbers 3) Fluorine has an oxidation number of 1 in all of its compounds because it is the most electronegative element. 4) Oxygen usually has an oxidation number of 2. Exceptions: i) In peroxides, such as H 2 O 2, oxygen s oxidation number is 1. ii) In compounds with fluorine, such as OF 2, oxygen s oxidation number is +2.
57 Assigning Oxidation Numbers 5) Hydrogen has an oxidation number of +1 in all compounds containing elements that are more electronegative than it; it has an oxidation number of 1 with metals. 6) The algebraic sum of the oxidation numbers of all atoms in an neutral compound is equal to zero. 7) The algebraic sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion. 8) Although rules 1 through 7 apply to covalently bonded atoms, oxidation numbers can also be applied to atoms in ionic compounds similarly.
58 Rules for Assigning Oxidation Numbers
59 Assigning Oxidation Numbers Assign oxidation numbers to each atom in the following compounds or ions: a. UF 6 b. H 2 SO 4 c. ClO 3 -
60 Assigning Oxidation Numbers a. UF 6 Place known oxidation numbers above the appropriate elements. 1 UF 6 Multiply known oxidation numbers by the appropriate number of atoms and place the totals underneath the corresponding elements. The compound UF 6 is molecular. The sum of the oxidation numbers must equal zero; therefore, the total of positive oxidation numbers is UF UF
61 Assigning Oxidation Numbers a. UF 6 Divide the total calculated oxidation number by the appropriate number of atoms. There is only one uranium atom in the molecule, so it must have an oxidation number of UF 6 UF UF UF
62 Assigning Oxidation Numbers b. H 2 SO 4 Hydrogen has an oxidation number of +1. Oxygen has an oxidation number of 2. The sum of the oxidation numbers must equal zero, and there is only one sulfur atom in each molecule of H 2 SO 4 Because (+2) + ( 8) = 6, the oxidation number of each sulfur atom must be H SO
63 Assigning Oxidation Numbers c. ClO 3 - The total of the oxidation numbers should equal the overall charge of the anion, 1. The oxidation number of a single oxygen atom in the ion is 2. The total oxidation number due to the three oxygen atoms is 6. For the chlorate ion to have a 1 charge, chlorine must be assigned an oxidation number of +5
64 Using Oxidation Numbers for Formulas and Names As shown in the table in the next slide, many nonmetals can have more than one oxidation number. These numbers can sometimes be used in the same manner as ionic charges to determine formulas. Example: What is the formula of a binary compound formed between sulfur and oxygen? From the common +4 and +6 oxidation states of sulfur, you could predict that sulfur might form SO 2 or SO 3 - Both are known compounds.
65 Common Oxidation States of Nonmetals
66 Using Oxidation Numbers for Formulas and Names Using oxidation numbers, the Stock system, introduced in the previous section for naming ionic compounds, can be used as an alternative to the prefix system for naming binary molecular compounds. PREFIX SYSTEM STOCK SYSTEM PCl 3 phosphorus trichloride phosphorus(iii) chloride PCl 5 phosphorus pentachloride phosphorus(v) chloride N 2 O dinitrogen monoxide nitrogen(i) oxide NO nitrogen monoxide nitrogen(ii) oxide Mo 2 O 3 dimolybdenum trioxide molybdenum(iii) oxide
67 Objectives Chapter 7, Section 7.3 Using Chemical Formulas i) Calculate the formula mass or molar mass of any given compound. ii) Use molar mass to convert between mass in grams and amount in moles of a chemical compound. iii) Calculate the number of molecules, formula units, or ions in a given molar amount of a chemical compound. iv) Calculate the percentage composition of a given chemical compound.
68 A chemical formula indicates: i) The elements present in a compound ii) The relative number of atoms or ions of each element present in a compound Chemical formulas also allow calculations of a number of other characteristic values for a compound: i) Formula mass ii) Molar mass iii) Percentage composition
69 Formula Masses The formula mass of any molecule, formula unit, or ion is the sum of the average atomic masses of all atoms represented in its formula. Example: Formula mass of water, H 2 O i) Average atomic mass of H: 1.01 amu ii) Average atomic mass of O: amu 1.01 amu 2 H atoms 2.02 amu H atom amu 1 O atom amu O atom Average mass of H 2 O molecule= = amu
70 Formula Masses The mass of a water molecule can be referred to as a molecular mass. The mass of one formula unit of an ionic compound, such as NaCl, is not a molecular mass. The mass of any unit represented by a chemical formula (H 2 O, NaCl) can be referred to as the formula mass.
71 Formula Mass
72 Formula Masses Sample Problem Find the formula mass of potassium chlorate, KClO 3 The mass of a formula unit of KClO 3 is found by adding the masses of one K atom, one Cl atom, and three O atoms. Atomic masses can be found in the periodic table in the back of your book. In your calculations, round each atomic mass to two decimal places.
73 Formula Masses amu 1 K atom amu K atom amu 1 Cl atom amu Cl atom amu 3 O atoms amu O atom Formula mass of KClO 3 = = amu
74 The Mole
75 Molar Masses The molar mass of a substance is equal to the mass in grams of one mole, or approximately particles, of the substance. Example: The molar mass of pure calcium, Ca, is g/mol because one mole of calcium atoms has a mass of g. The molar mass of a compound is calculated by adding the masses of the elements present in a mole of the molecules or formula units that make up the compound.
76 Molar Masses One mole of water molecules contains exactly two moles of H atoms and one mole of O atoms. The molar mass of water is calculated as follows g H 2 mol H 2.02 g H mol H g O 1 mol O g O mol O molar mass of H 2 O molecule: g/mol A compound s molar mass is numerically equal to its formula mass.
77 Calculating Molar Masses of Ionic Compounds
78 Molar Masses Sample Problem What is the molar mass of barium nitrate, Ba(NO 3 ) 2? One mole of barium nitrate, contains one mole of Ba, two moles of N (1 2), and six moles of O (3 2) g H 1 mol Ba g Ba mol Ba g 2 mol N g N mol N g O 6 mol O g O mol O molar mass of Ba(NO 3 ) 2 = = g/mol
79 Molar Mass as a Conversion Factor The molar mass of a compound can be used as a conversion factor to relate an amount in moles to a mass in grams for a given substance. To convert moles to grams, multiply the amount in moles by the molar mass: Amount in moles molar mass (g/mol) = mass in grams
80 Mole-Mass Calculations
81 Molar Mass as a Conversion Factor
82 Molar Mass as a Conversion Factor Sample Problem What is the mass in grams of 2.50 mol of oxygen gas? Given: 2.50 mol O 2 Unknown: mass of O 2 in grams Solution: moles O 2 grams O2 amount of O 2 (mol) molar mass of O 2 (g/mol) = mass of O 2 (g)
83 Molar Mass as a Conversion Factor Sample Problem Calculate the molar mass of O2: g O 2 mol O g mol O Use the molar mass of O 2 to convert moles to mass g O mol O mol O2 go g O 2
84 Converting Between Amount in Moles and Number of Particles
85 Molar Mass as a Conversion Factor Sample Problem Ibuprofen, C 13 H 18 O 2, is the active ingredient in many nonprescription pain relievers. Its molar mass is g/mol. i) If the tablets in a bottle contain a total of 33 g of ibuprofen, how many moles of ibuprofen are in the bottle? ii) How many molecules of ibuprofen are in the bottle? iii) What is the total mass in grams of carbon in 33 g of ibuprofen?
86 Molar Mass as a Conversion Factor Solution - i) If the tablets in a bottle contain a total of 33 g of ibuprofen, how many moles of ibuprofen are in the bottle? Given: 33 g of C 13 H 18 O 2 molar mass g/mol Unknown: a. moles C 13 H 18 O 2 b. molecules C 13 H 18 O 2 c. total mass of C g C H O i) grams moles 1 mol C H O mol C H O g C13H 18O2
87 Molar Mass as a Conversion Factor Solution - ii) How many molecules of ibuprofen are in the bottle? moles molecules molecules molecules C13H 18O2 mol C H O mol Solution - iii) What is the total mass in grams of carbon in 33 g of ibuprofen? moles C 13 H 18 O 2 moles C grams C mol C H O mol C g C mol C H O mol C g C
88 Molar Mass as a Conversion Factor Solution - i) If the tablets in a bottle contain a total of 33 g of ibuprofen, how many moles of ibuprofen are in the bottle? 33 g C H O mol C H O g C H O mol C H O Solution - ii) How many molecules of ibuprofen are in the bottle? 0.16mol C H O Solution - iii) What is the total mass in grams of carbon in 33 g of ibuprofen? molecules mol molecules C13H18O2 13 mol C g C 0.16 mol C H O 25 g C mol C H O mol C
89 Percentage Composition It is often useful to know the percentage by mass of a particular element in a chemical compound. To find the mass percentage of an element in a compound, the following equation can be used. mass of element in sample of compound 100 mass of sample of compound % element in compound The mass percentage of an element in a compound is the same regardless of the sample s size.
90 Percentage Composition The percentage of an element in a compound can be calculated by determining how many grams of the element are present in one mole of the compound. mass of element in 1 mol of compound 100 molar mass of compound % element in compound The percentage by mass of each element in a compound is known as the percentage composition of the compound.
91 Percentage Composition of Iron Oxides
92 Percentage Composition
93 Percentage Composition Calculations
94 Percentage Composition Sample Problem Find the percentage composition of copper(i) sulfide, Cu 2 S. Given: Unknown: Formula, Cu 2 S Percentage composition of Cu 2 S Solution: formula molar mass mass percentage of each element g Cu 2 mol Cu g Cu mol Cu g S 1 mol S g S mol S Molar mass of Cu 2 S = = g
95 Percentage Composition Solution (Continued) g Cu g Cu S % Cu g S g CuS % S
96 Objectives Chapter 7, Section 7.4 Determining Chemical Formulas i) Define empirical formula, and explain how the term applies to ionic and molecular compounds. ii) Determine an empirical formula from either a percentage or a mass composition. iii) Explain the relationship between the empirical formula and the molecular formula of a given compound. iv) Determine a molecular formula from an empirical formula.
97 Empirical and Actual Formulas
98 Empirical and Actual Formulas An empirical formula consists of the symbols for the elements combined in a compound, with subscripts showing the smallest wholenumber 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. Example: the empirical formula of the gas diborane is BH 3, but the molecular formula is B 2 H 6.
99 Calculation of Empirical Formulas To determine a compound s empirical formula from its percentage composition, begin by converting percentage composition to a mass composition. Assume that you have a g sample of the compound. Then calculate the amount of each element in the sample. Example: Diborane, BH 3 The percentage composition is 78.1% B and 21.9% H. Therefore, g of diborane contains 78.1 g of B and 21.9 g of H.
100 Calculation of Empirical Formulas Next, the mass composition of each element is converted to a composition in moles by dividing by the appropriate molar mass. 1 mol B 78.1 g B 7.22 mol B g B 1 mol H 21.9 g H 21.7 mol H 1.01 g H These values give a mole ratio of 7.22 mol B to 21.7 mol H. To find the smallest whole number ratio, divide each number of moles by the smallest number in the existing ratio mol B 21.7 mol H : 1 mol B : 3.01 mol H
101 Calculation of Empirical Formulas Because of rounding or experimental error, a compound s mole ratio sometimes consists of numbers close to whole numbers instead of exact whole numbers. In this case, the differences from whole numbers may be ignored and the nearest whole number taken.
102 Percentage Composition Sample Problem Quantitative analysis shows that a compound contains 32.38% sodium, 22.65% sulfur, and 44.99% oxygen. Find the empirical formula of this compound. Solution Given: Percentage composition: 32.38% Na, 22.65% S, and 44.99% O Unknown: Empirical formula percentage composition mass composition composition in moles smallest whole-number mole ratio of atoms
103 Percentage Composition Sample Problem Solution 1 mol Na g Na mol Na g Na 1 mol S g S mol S g S 1 mol O g O mol O g O Smallest whole-number mole ratio of atoms: The compound contains atoms in the ratio mol Na; and mol S; and mol O mol Na mol S mol O : : mol Na :1 mol S : mol O
104 Percentage Composition Sample Problem Solution Rounding yields a mole ratio of: i) 2 mol Na ii) 1 mol S iii) 4 mol O The empirical formula of the compound is Na 2 SO 4
105 Calculation of Molecular Formulas The empirical formula contains the smallest possible whole numbers that describe the atomic ratio. The molecular formula is the actual formula of a molecular compound. An empirical formula may or may not be a correct molecular formula. The relationship between a compound s empirical formula and its molecular formula can be written as follows. (X)(Empirical Formula) = (Molecular Formula)
106 Calculation of Molecular Formulas The formula masses have a similar relationship. (X)(Empirical Formula Mass) = (Molecular Formula Mass) To determine the molecular formula of a compound, you must know the compound s formula mass. Dividing the experimental formula mass by the empirical formula mass gives the value of X A compound s molecular formula mass is numerically equal to its molar mass, so a compound s molecular formula can also be found given the compound s empirical formula and its molar mass.
107 Comparing Empirical and Molecular Formulas
108 Comparing Molecular and Empirical Formulas
109 Percentage Composition Sample Problem In Sample Problem M, the empirical formula of a compound of phosphorus and oxygen was found to be P 2 O 5. Experimentation shows that the molar mass of this compound is g/mol. What is the compound s molecular formula? Given: Unknown: Solution: Empirical Formula Molecular Formula (X)(Empirical Formula) = (Molecular Formula) x molecular formula mass empirical formula mass
110 Percentage Composition Sample Problem Solution, continued Molecular formula mass is numerically equal to molar mass. i) Molecular Molar Mass = g/mol ii) Molecular Formula Mass = amu Empirical Formula Mass i) mass of phosphorus atom = amu ii) mass of oxygen atom = amu Empirical Formula Mass of P 2 O 5 = amu amu = amu
111 Percentage Composition Sample Problem Solution, continued Dividing the experimental formula mass by the empirical formula mass gives the value of x. x = amu amu 2 (P 2 O 5 ) = P 4 O 10 The compound s molecular formula is therefore P 4 O 10
SCH 4C1 Unit 2 Problem Set Questions taken from Frank Mustoe et all, "Chemistry 11", McGraw-Hill Ryerson, 2001 1. A small pin contains 0.0178 mol of iron. How many atoms of iron are in the pin? 2. A sample
1 MOLECULAR MASS AND FORMULA MASS Molecular mass = sum of the atomic weights of all atoms in the molecule. Formula mass = sum of the atomic weights of all atoms in the formula unit. 2 MOLECULAR MASS AND
35 MOLES ND MOLE CLCULTIONS INTRODUCTION The purpose of this section is to present some methods for calculating both how much of each reactant is used in a chemical reaction, and how much of each product
Chapter 3 Mass Relationships in Chemical Reactions Student: 1. An atom of bromine has a mass about four times greater than that of an atom of neon. Which choice makes the correct comparison of the relative
Chapter 3: Stoichiometry Key Skills: Balance chemical equations Predict the products of simple combination, decomposition, and combustion reactions. Calculate formula weights Convert grams to moles and
Ch. 10 The Mole I. Molar Conversions I II III IV A. What is the Mole? A counting number (like a dozen) Avogadro s number (N A ) 1 mole = 6.022 10 23 representative particles B. Mole/Particle Conversions
Chapter 3. Stoichiometry: Mole-Mass Relationships in Chemical Reactions Concept 1. The meaning and usefulness of the mole The mole (or mol) represents a certain number of objects. SI def.: the amount of
CHEMISTRY 123-07 Midterm #1 Answer key October 14, 2010 Statistics: Average: 74 p (74%); Highest: 97 p (95%); Lowest: 33 p (33%) Number of students performing at or above average: 67 (57%) Number of students
How much does a single atom weigh? Different elements weigh different amounts related to what makes them unique. What units do we use to define the weight of an atom? amu units of atomic weight. (atomic
CHAPTER 3 Calculations with Chemical Formulas and Equations MOLECULAR WEIGHT (M. W.) Sum of the Atomic Weights of all atoms in a MOLECULE of a substance. FORMULA WEIGHT (F. W.) Sum of the atomic Weights
Chapter 1: Moles and equations 1 Learning outcomes you should be able to: define and use the terms: relative atomic mass, isotopic mass and formula mass based on the 12 C scale perform calculations, including
Pre- AP Chemistry Chemical Quan44es: The Mole Mole SI unit of measurement that measures the amount of substance. A substance exists as representa9ve par9cles. Representa9ve par9cles can be atoms, molecules,
Moles, Molecules, and Grams Worksheet Answer Key 1) How many are there in 24 grams of FeF 3? 1.28 x 10 23 2) How many are there in 450 grams of Na 2 SO 4? 1.91 x 10 24 3) How many grams are there in 2.3
Using Chemical Formulas Objective 1: Calculate the formula mass or molar mass of any given compound. The Formula Mass of any molecule, formula unit, or ion is the sum of the average atomic masses of all
10 CHEMICAL QUANTITIES SECTION 10.1 THE MOLE: A MEASUREMENT OF MATTER (pages 287 296) This section defines the mole and explains how the mole is used to measure matter. It also teaches you how to calculate
Molar Mass Molar mass = Mass in grams of one mole of any element, numerically equal to its atomic weight Molar mass of molecules can be determined from the chemical formula and molar masses of elements
Chapter 3 Stoichiometry 3-1 Chapter 3 Stoichiometry In This Chapter As you have learned in previous chapters, much of chemistry involves using macroscopic measurements to deduce what happens between atoms
TEKS REVIEW 8B Calculating Atoms, Ions, or Molecules Using Moles TEKS 8B READINESS Use the mole concept to calculate the number of atoms, ions, or molecules in a sample TEKS_TXT of material. Vocabulary
Chemical Composition Review Mole Calculations Percent Composition Copyright Cengage Learning. All rights reserved. 8 1 QUESTION Suppose you work in a hardware store and a customer wants to purchase 500
T-6 Tutorial 2 FORMULAS, PERCENTAGE COMPOSITION, AND THE MOLE FORMULAS: A chemical formula shows the elemental composition of a substance: the chemical symbols show what elements are present and the numerical
1. When the following equation is balanced, the coefficient of Al is. Al (s) + H 2 O (l)? Al(OH) (s) + H 2 (g) A) 1 B) 2 C) 4 D) 5 E) Al (s) + H 2 O (l)? Al(OH) (s) + H 2 (g) Al (s) + H 2 O (l)? Al(OH)
Moles Balanced chemical equations Molar ratios Mass Composition Empirical and Molecular Mass Predicting Quantities Equations Micro World atoms & molecules Macro World grams Atomic mass is the mass of an
Name: Chemistry Post-Enrolment Worksheet The purpose of this worksheet is to get you to recap some of the fundamental concepts that you studied at GCSE and introduce some of the concepts that will be part
Chapter 3 ATOMS AND MOLECULES Multiple Choice Questions 1. Which of the following correctly represents 360 g of water? (i) 2 moles of H 2 0 (ii) 20 moles of water (iii) 6.022 10 23 molecules of water (iv)
Chapter 1 The Atomic Nature of Matter 6. Substances that cannot be decomposed into two or more simpler substances by chemical means are called a. pure substances. b. compounds. c. molecules. d. elements.
Example Exercise 9.1 Atomic Mass and Avogadro s Number Refer to the atomic masses in the periodic table inside the front cover of this textbook. State the mass of Avogadro s number of atoms for each of
How do we figure this out? We know that: 1) the number of oxygen atoms can be found by using Avogadro s number, if we know the moles of oxygen atoms; 2) the number of moles of oxygen atoms can be found
LOCUS 1 Section - CHEMICAL FORMULAE AND STOICHIOMETRY The language that chemists use to describe the forms of matter and the changes in its composition appears throughout the scientific world. Chemical
Chapter 3 Formulas, Equations and Moles Interpreting Chemical Equations You can interpret a balanced chemical equation in many ways. On a microscopic level, two molecules of H 2 react with one molecule
Woods Chem-1 Lec-02 10-1 Atoms, Ions, Mole (std) Page 1 ATOMIC THEORY, MOLECULES, & IONS Proton: A positively charged particle in the nucleus Atomic Number: We differentiate all elements by their number
Sample Exercise 3.1 Interpreting and Balancing Chemical Equations The following diagram represents a chemical reaction in which the red spheres are oxygen atoms and the blue spheres are nitrogen atoms.
STOICHIOMETRY LEARNING OUTCOMES At the end of this unit students will be expected to: UNIT 1 THE MOLE AND MOLAR MASS define molar mass and perform mole-mass inter-conversions for pure substances explain
Ch. 6 Chemical Composition and Stoichiometry The Mole Concept [6.2, 6.3] Conversions between g mol atoms [6.3, 6.4, 6.5] Mass Percent [6.6, 6.7] Empirical and Molecular Formula [6.8, 6.9] Bring your calculators!
The Mole Concept Ron Robertson r2 c:\files\courses\1110-20\2010 final slides for web\mole concept.docx The Mole The mole is a unit of measurement equal to 6.022 x 10 23 things (to 4 sf) just like there
The Mole Chapter 10 1 Objectives Use the mole and molar mass to make conversions among moles, mass, and number of particles Determine the percent composition of the components of a compound Calculate empirical
Chem 31 Fall 2002 Chapter 3 Stoichiometry: Calculations with Chemical Formulas and Equations Writing and Balancing Chemical Equations 1. Write Equation in Words -you cannot write an equation unless you
Micro World atoms & molecules Laboratory scale measurements Atomic mass is the mass of an atom in atomic mass units (amu) By definition: 1 atom 12 C weighs 12 amu On this scale 1 H = 1.008 amu 16 O = 16.00
CHAPTER THREE: CALCULATIONS WITH CHEMICAL FORMULAS AND EQUATIONS Part One: Mass and Moles of Substance A. Molecular Mass and Formula Mass. (Section 3.1) 1. Just as we can talk about mass of one atom of
Name: Class: Date: Chemistry Final Study Guide Multiple Choice Identify the choice that best completes the statement or answers the question. 1. The electrons involved in the formation of a covalent bond
General Chemistry Principles and Modern Applications Petrucci Harwood Herring 8 th Edition Chapter 3: Chemical Compounds Dr. Burak Esat Fatih University Chem 107Fall 2013 1 Compound: A combination of two
Lecture 5, The Mole What is a mole? Moles Atomic mass unit and the mole amu definition: 12 C = 12 amu. The atomic mass unit is defined this way. 1 amu = 1.6605 x 10-24 g How many 12 C atoms weigh 12 g?
EXPERIMENT 12: Empirical Formula of a Compound INTRODUCTION Chemical formulas indicate the composition of compounds. A formula that gives only the simplest ratio of the relative number of atoms in a compound
Contents Getting the most from this book...4 About this book....5 Content Guidance Topic 1 Atomic structure and the periodic table...8 Topic 2 Bonding and structure...14 Topic 2A Bonding....14 Topic 2B
Instructions for Conversion Problems For every conversion problem Write the number in the problem down with unit and a multiplication sign Decide which conversion factor you should use, Avagadro s or molar
Matter Atomic weight, Molecular weight and Mole Atomic Mass Unit Chemists of the nineteenth century realized that, in order to measure the mass of an atomic particle, it was useless to use the standard
Candidate Style Answer Chemistry A Unit F321 Atoms, Bonds and Groups High banded response This Support Material booklet is designed to accompany the OCR GCE Chemistry A Specimen Paper F321 for teaching
29 Chemical Formulae Chemical formulae are used as shorthand to indicate how many atoms of one element combine with another element to form a compound. C 2 H 6, 2 atoms of carbon combine with 6 atoms of
Calculations and Chemical Equations Atomic mass: Mass of an atom of an element, expressed in atomic mass units Atomic mass unit (amu): 1.661 x 10-24 g Atomic weight: Average mass of all isotopes of a given
96 Chapter 7: Calculations with Chemical Formulas and Chemical Reactions Chemical reactions are written showing a few individual atoms or molecules reacting to form a few atoms or molecules of products.
A dozen Molar Mass Science 10 is a number of objects. A dozen eggs, a dozen cars, and a dozen people are all 12 objects. But a dozen cars has a much greater mass than a dozen eggs because the mass of each
Test Review Periodic Trends and The Mole The Mole SHOW ALL WORK ON YOUR OWN PAPER FOR CREDIT!! 1 2 (NH42SO2 %N 24.1 %H 6.9 %S 27.6 %O 41.3 % Al %C 35.3 %H 4.4 %O 47.1 Al(C2H3O23 13.2 3 How many moles are
CHEM 120 Online: Chapter 6 Sample problems Date: 1. To determine the formula mass of a compound you should A) add up the atomic masses of all the atoms present. B) add up the atomic masses of all the atoms
Tuesday, November 27, 2012 Expectations: Sit in assigned seat Get out Folder, Notebook, Periodic Table Have out: Spiral (notes), Learning Target Log (new) No Backpacks on tables Listen/Pay Attention Learning
3.3 Moles, 3.4 Molar Mass, and 3.5 Percent Composition Collection Terms A collection term states a specific number of items. 1 dozen donuts = 12 donuts 1 ream of paper = 500 sheets 1 case = 24 cans Copyright
AP Chapter 1, 2, & 3: Atoms, Molecules, and Mass Relationships Name Warm-Ups (Show your work for credit) Date 1. Date 2. Date 3. Date 4. Date 5. Date 6. Date 7. Date 8. AP Chapter 1, 2, & 3: Atoms & Molecules,
Atomic Masses Chapter 3 Stoichiometry 1 atomic mass unit (amu) = 1/12 of the mass of a 12 C atom so one 12 C atom has a mass of 12 amu (exact number). From mass spectrometry: 13 C/ 12 C = 1.0836129 amu
Mole Calculations Chemical Equations and Stoichiometry Lecture Topics Atomic weight, Mole, Molecular Mass, Derivation of Formulas, Percent Composition Chemical Equations and Problems Based on Miscellaneous
44 Section 43 Questions 1 Define Avogadro s constant, and explain its significance in quantitative analysis 2 Distinguish between the terms atomic mass and molar mass 3 Calculate the mass of a molecule
sil07204_ch03_69-107 8/22/05 15:16 Page 69 CHAPTER THREE Stoichiometry of Formulas and Equations Key Principles The mole (mol) is the standard unit for amount of substance and consists of Avogadro s number
Chapter 5 Chemical Quantities and Reactions 5.1 The Mole Collection Terms A collection term states a specific number of items. 1 dozen donuts = 12 donuts 1 ream of paper = 500 sheets 1 case = 24 cans 1
Chapter 7 Answers and Solutions 7 Answers and Solutions to Text Problems 7.1 A mole is the amount of a substance that contains 6.02 x 10 23 items. For example, one mole of water contains 6.02 10 23 molecules
7 Chemical Composition 7. Avogadro s umber n 8, the talian scientist, Amadeo Avogadro proposed that: Equal volumes of gas at equal temperatures and pressures have the same number of particles. This law,
MASS RELATIONSHIPS IN CHEMICAL REACTIONS 1. The mole, Avogadro s number and molar mass of an element. Molecular mass (molecular weight) 3. Percent composition of compounds 4. Empirical and Molecular formulas
SCH3U- R.H.KING ACADEMY SOLUTION & ACID/BASE WORKSHEET Name: The importance of water - MAKING CONNECTION READING 1. Read P. 368-375, P. 382-387 & P. 429-436; P. 375 # 1-11 & P. 389 # 1,7,9,12,15; P. 436
Bonding in Elements and Compounds Structure of solids, liquids and gases Types of bonding between atoms and molecules Ionic Covalent Metallic Many compounds between metals & nonmetals (salts), e.g. Na,
hapter 2: Atoms and Molecules HAPTER OUTLINE 2.1 Symbols and Formulas 2.2 Inside the Atom 2. Isotopes 2.4 Relative Masses of Atoms and Molecules 2.5 Isotopes and Atomic Weights 2.6 Avogadro s Number: The
CHM111(M)/Page 1 of 5 INTI COLLEGE MALAYSIA A? LEVEL PROGRAMME CHM 111: CHEMISTRY MOCK EXAMINATION: DECEMBER 2000 SESSION SECTION A Answer ALL EIGHT questions. (52 marks) 1. The following is the mass spectrum
Moles I. Lab: Rice Counting II. Counting atoms and molecules I. When doing reactions chemists need to count atoms and molecules. The problem of actually counting individual atoms and molecules comes from
CHE 100 Worksheet # 8 Graham/09 Name Key Due 1. According to the law of definite proportions, if a sample of a compound contains 7.00 grams of sulfur and 3.50 grams of oxygen, then another sample of the
Lake County, Lakeview, 9 th grade, Physical Science, Brent Starr Standard: H1P1 Explain how atomic structure is related to the properties of elements and their position in the Periodic Table. Explain how
Name The Mole Lab Chemistry I Acc (Weihin as a Means of Countin) Introduction Date One of the seven SI base units is the mole. The mole, also known as Avoadro s number, is equal to 6.02 x 10 23. The mole
Name: Class: Date: Chapter 13 & 14 Practice Exam Multiple Choice Identify the choice that best completes the statement or answers the question. 1. Acids generally release H 2 gas when they react with a.
Chapter 5 s to Supplementary Check for Understanding Problems Moles and Molar Mass 1. Indicate the appropriate quantity for each of the following. a) A mole of N atoms contains atoms. b) A mole of N molecules
LAB 5 PLANT NUTRITION I. General Introduction All living organisms require certain elements for their survival. Plants are known to require carbon (C), hydrogen (H), oxygen (O), nitrogen (N), phosphorus
Bonds hapter 8 Bonding: General oncepts Forces that hold groups of atoms together and make them function as a unit. Bond Energy Bond Length It is the energy required to break a bond. The distance where
Atoms, Ions and Molecules The Building Blocks of Matter Chapter 2 1 Chapter Outline 2.1 The Rutherford Model of Atomic Structure 2.2 Nuclides and Their Symbols 2.3 Navigating the Periodic Table 2.4 The