Chapter 2 Elements, Compounds, and the Periodic Table

Chapter 2 Elements, Compounds, and the Periodic Table
Chemistry: The Molecular Nature of Matter, 7E Jespersen/Hyslop

Chapter in Context Explore the distribution of metals, nonmetals, and metalloids in the periodic table Interpret chemical formulas Understand balanced chemical equations and how they relate to atomic theory Use the periodic table and ion charges to write formulas for ionic compounds Name ionic formulas and write formulas from chemical names Understand the difference between ionic and molecular compounds Name molecular compounds

Periodic Table Early Versions of Periodic Tables Modern Periodic Table
Summarizes periodic properties of elements Early Versions of Periodic Tables Arranged by increasing atomic mass Mendeleev (Russian) and Meyer (German) in 1869 Noted repeating (periodic) properties Modern Periodic Table Arranged by increasing atomic number (Z ): Rows called periods Columns called groups or families Identified by numbers 1 – 18 standard international 1A – 8A longer columns and 1B – 8B shorter columns 3 3

Modern Periodic Table with group labels and chemical families identified Note: Placement of elements 58 – 71 and 90 – 103 saves space 4 4

Representative/Main Group Elements
A groups—Longer columns Alkali Metals 1A = first group Very reactive All are metals except for H Tend to form +1 ions React with oxygen Form compounds that dissolve in water Yield strongly caustic or alkaline solution (Na2O) 5 5

A groups—Longer columns Alkaline Earth Metals 2A = second group Reactive Tend to form +2 ions Oxygen compounds are strongly alkaline (MgO) Many are not water soluble 6

A groups—Longer columns Halogens 7A = next to last group on right Reactive Form diatomic molecules in elemental state 2 gases – F2, Cl2 1 liquid – Br2 2 solids – I2, At2 Form –1 ions with alkali metals—salts (e.g. NaF, NaCl, NaBr, and NaI) 7

A groups—Longer columns Noble Gases 8A = last group on right Inert—very unreactive Only heavier elements of group react and then very limited Don’t form charged ions Monatomic gases (e.g., He, Ne, Ar) 8

Transition Elements B groups—shorter columns e.g., All are metals
In center of table Begin in fourth row Tend to form ions with several different charges e.g., Fe2+ and Fe3+ Cu+ and Cu2+ Mn2+, Mn3+, Mn4+, Mn5+, Mn6+, and Mn7+ Note: Last 3 columns all have 8B designation 9 9

Inner Transition Elements
At bottom of periodic table Tend to form +2 and +3 ions Lanthanide elements Elements 58 – 71 Actinide elements Elements 90 – 103 All actinides are radioactive 10 10

Metals, Nonmetals, or Metalloids
Elements break down into three broad categories Organized by regions of periodic table Metals Left-hand side Sodium, lead, iron, gold Nonmetals Upper right-hand corner Oxygen, nitrogen, chlorine Metalloids Diagonal line between metals and nonmetals Boron to astatine 11 11

Metals, Nonmetals, or Metalloids
12 12

Your Turn! Classify the following three elements as a metal, non-metal, or metalloid: silicon (Si), vanadium (V), bromine (Br) nonmetal, metal, nonmetal, respectively metal, metalloid, nonmetal, respectively nonmetal, metal, metalloid, respectively metalloid, metal, metalloid, respectively None of these are correct 13

Your Turn! Strontium (Sr) is a _______, ruthenium (Ru) is a ________, and iodine (I) is a_________. alkali metal, transition metal, halogen transition metal, alkaline earth metal, halogen alkaline earth metal, transition metal, halogen transition metal, alkali metal, noble gas alkali metal, actinide, halogen 14

Metals Properties Most elements in periodic table Metallic luster
Shine or reflect light Malleable Can be hammered or rolled into thin sheets Ductile Can be drawn into wire Hardness Some hard – iron and chromium Some soft – sodium, lead, copper 15

Properties of Metals Conduct heat and electricity
Solids at room temperature Melting points (mp) > 25 °C Hg only liquid metal (mp = –39 °C) Tungsten (W) (mp = 3400 °C) Highest mp for a metal Chemical reactivity Varies greatly Au, Pt very unreactive Na, K very reactive 16 16

Nonmetals Seventeen elements
Upper right hand corner of periodic table Exist mostly as compounds rather than as pure elements Many are gases Monatomic (Noble) He, Ne, Ar, Kr, Xe, Rn Diatomic H2, O2, N2, F2, Cl2 Some are solids: I2, Se8, S8, P4, C Three forms of carbon (graphite, coal, diamond) One is liquid: Br2 17 17

Properties of Nonmetals
Brittle Pulverize when struck Insulators Non-conductors of electricity and heat Chemical reactivity Some inert Noble gases Some reactive F2, O2, H2 React with metals to form ionic compounds 18 18

Metalloids Eight Elements Properties Semiconductors
Located on diagonal line between metals and nonmetals B, Si, Ge, As, Sb, Te, Po, At Properties Between metals and nonmetals Metallic shine Brittle like nonmetal Semiconductors Conduct electricity But not as well as metals Silicon (Si) and germanium (Ge) 19 19

Your Turn! Which of the following statements is correct?
Cu is a representative transition element Na is an alkaline earth metal Al is a metalloid in group 3A F is a representative halogen None of these are correct 20

Your Turn! All of the following are characteristics of metals except:
Malleable Ductile Lustrous Good conductors of heat Act as semiconductors 21

Your Turn! Nonmetals such as bromine and oxygen: Conduct heat
React with metals Are lustrous Conduct electricity Acts as semiconductors 22

Molecules and Chemical Formulas
Atoms combine into compounds Useful to visualize atoms, compounds, and molecules Atoms represented by spheres Different atoms have different colors Standard scheme is represented on the right

Molecules Atoms combine to form more complex substances
Discrete particles Each composed of two or more atoms e.g., Molecular oxygen, O2 Carbon dioxide, CO2 Ammonia, NH3 Sucrose, C12H22O11

Chemical Formulas Specify composition of substance Chemical symbols
Represent atoms of elements present Subscripts Given after chemical symbol Represents relative numbers of each type of atom e.g., Fe2O3 : iron and oxygen in 2:3 ratio

Chemical Formulas Free Elements e.g., Iron Fe Neon Ne
Element not combined with another in compounds Just use chemical symbol to represent e.g., Iron Fe Neon Ne Sodium Na Aluminum Al Diatomic Molecule Molecules composed of two atoms each Many elements found in nature e.g., Oxygen O2 Nitrogen N2 Hydrogen H2 Chlorine Cl2

Depicting Molecules Want to show: Three ways of visualizing molecules:
Order in which atoms are attached to each other 3-dimensional shape of molecule Three ways of visualizing molecules: Structural formula Ball-and-stick model Space filling model

1. Structural Formulas Use to show how atoms are attached
Atoms represented by chemical symbols Chemical bonds attaching atoms indicated by lines

3-D Representations of Molecules
Hydrogen molecule, H2 Oxygen molecule, O2 Nitrogen molecule N2 Chlorine molecule, Cl2 Use fused spheres to indicate molecules Different colors indicate different elements Relative size of spheres reflects differing sizes of atoms

2. “Ball-and-Stick” Model
Spheres = atoms Sticks = bonds Methane CH4 Chloroform CHCl3

3. “Space-Filling” Model
Shows relative sizes of atoms Shows how atoms take up space in molecule Water H2O Methane CH4 Chloroform, CHCl3

More Complicated Molecules
Sometimes formulas contain parentheses How do we translate into a structure? e.g., Urea, CO(NH2)2 Expands to CON2H4 Atoms in parentheses appear twice Ball-and-stick model Space-filling model

Hydrates Crystals that contain water molecules Dehydration
e.g., Plaster: CaSO4∙2H2O calcium sulfate dihydrate Water is not tightly held Dehydration Removal of water by heating Remaining solid is anhydrous (without water) Blue = CuSO4 •5H2O White = CuSO4

Counting Atoms Subscript following chemical symbol indicates how many of that element are part of the formula No subscript implies a subscript of 1 Quantity in parentheses is repeated a number of times equal to the subscript that follows Raised dot in formula indicates that the substance is a hydrate Number preceding H2O specifies how many water molecules are present

Counting Atoms Example: 1 (CH3)3COH Subscript 3 means 3 CH3 groups
So from (CH3)3 we get 3 × 1 C = 3 C 3 × 3 H = 9 H Number of C = 3 C + 1 C = 4 C Number of H = 9 H + 1 H = 10 H Number of O = 1 O Total number of atoms = 15 atoms

Counting Atoms Example: 2 CoCl2·6H2O
The dot 6H2O means you multiple both H2 and O by 6 So there are: Number of H 6 × 2 = 12 H Number of O 6 × 1 = 6 O Number of Co 1 × 1 = 1 Co Number of Cl 2 × 1 = 2 Cl Total Number of atoms = 21 atoms

Your Turn! Count the number of each type of atom in the chemical formula given below Na2CO3 (NH4)2SO4 Mg3(PO4)2 CuSO4∙5H2O (C2H5)2N2H2 ___Na, ___ C, ___ O ___N, ___H, ___S, ___O ___Mg, ___P, ___O ___Cu, ___S, ___O, ___H ___C, ___H, ___N 2 1 3 2 8 1 4 3 2 8 1 1 9 10 4 12 2

Your Turn! What are the formulas of the compounds shown?
C3H8, CH3NH2, C3H6O2 C3H7, CH5Cl, CH6O2 C3H8, CH3NH2, C3H6Br2 C3H8, CH3SiH2, C3H6O2 N3H8, CH3BH2, C3H6O2 38

Dalton’s Atomic Theory
We now have the tools to explain this theory and its consequences All molecules of compound are alike and contain atoms in same numerical ratio. Example: Water, H2O Ratio of oxygen to hydrogen is 1 : 2 1 O atom : 2 H atoms in each molecule O weighs 16 times as much as H 1 H = 1 mass unit 1 O = 16 mass units

Atoms in Fixed Ratios by Mass
For water in general: Mass O = 8 times the mass H Regardless of amount of water present

Dalton’s Atomic Theory
Successes: Explains Law of Conservation of Mass Chemical reactions correspond to rearranging atoms. Explains Law of Definite Proportions Given compound always has atoms of same elements in same ratios. Predicted Law of Multiple Proportions Not yet discovered Some elements combine to give two or more compounds e.g., SO2 and SO3

Law of Multiple Proportions
When two elements form more than one compound, different masses of one element that combine with same mass of other element are always in ratio of small whole numbers. Atoms react as complete (whole) particles. Chemical formulas Indicate whole numbers of atoms Not fractions

Using Law Of Multiple Proportions
sulfur sulfur dioxide trioxide Mass S g g Mass O g g Use this data to prove law of multiple proportions

Law of Multiple Proportions
Compound Sample Size Mass of Sulfur Mass of Oxygen Sulfur dioxide 64.06 g 32.06 g 32.00 g Sulfur trioxide 80.06 g 32.06 g 48.00 g Ratio of

Molecules Small and Large
So far we’ve only discussed small molecules Some are very large, especially those found in nature Same principles apply to all e.g., DNA - short segment

How Do We Know Formulas? Hardly “out of the blue”
Don’t know formula when compound first isolated Formulas and structures backed by extensive experimentation Use results of experiments to determine Formula Chemical reactivity Molecular shape Can speculate once formula is known Determine from more experiments

Visualizing Mixtures Look at mixtures at atomic/molecular level
Different color spheres stand for two substances Homogeneous mixture/solution – uniform mixing Heterogeneous mixture – two phases a. b.

Chemical Reactions When one or more substances react to form one or more new substances e.g., Reaction of methane, CH4, with oxygen, O2, to form carbon dioxide, CO2, and water, H2O. Reactants = CH4 and O2 Products = CO2 and H2O How to depict? Words too long Pictures too awkward

Chemical Equations Use chemical symbols and formulas to represent reactants and products. Reactants on left hand side Products on right hand side Arrow () means “reacts to yield” e.g., CH4 + 2O2  CO2 + 2H2O Coefficients Numbers in front of formulas Indicate how many of each type of molecule reacted or formed Equation reads “methane and oxygen react to yield carbon dioxide and water”

Conservation of Mass in Reactions
Mass can neither be created nor destroyed This means that there are the same number of each type of atom in reactants and in products of reaction If number of atoms same, then mass also same CH4 + 2O2  CO2 + 2H2O 4H + 4O + C = 4H + 4O + C

Balanced Chemical Equation
Ex. 2C4H O2  8CO2 + 10H2O 4 C and 10 H per molecule 2 O per molecule 2 H and 1 O per molecule 1 C and 2 O per molecule Subscripts Define identity of substances Must not change when equation is balanced

Balanced Chemical Equation
Ex. 2C4H O2  8CO2 + 10H2O 8 molecules of CO2 2 molecules of C4H10 13 molecules of O2 10 molecules of C4H10 Coefficients Number in front of formulas Indicate number of molecules of each type Adjusted so number of each type of atom is same on both sides of arrow Can change

Balanced Chemical Equations
How do you determine if an equation is balanced? Count atoms Same number of each type on both sides of equation? If yes, then balanced If no, then unbalanced Ex. 2C4H O2  8CO2 + 10H2O Reactants Products 2×4 = 8 C 8×1 = 8 C 2×10 = 20 H 10×2 = 20 H 13×2 = 26 O (8×2)+(10×1)= 26 O

Learning Check Fe(OH)3 + 2 HNO3  Fe(NO3)3 + 2 H2O Not balanced
Only Fe has same number of atoms on either side of arrow. Reactants Products Fe 1 3 + (2×3) = 9 (3×3) + 2 = 11 O 3 + 2 = 5 (2×2) = 4 H 2 3 N

Your Turn! How many atoms of each element appear on each side of the arrow in the following equation? 4NH3 + 3O2 → 2N2 + 6H2O Reactants Products N (4 × 1) = 4 (2 × 2) = 4 O (3 × 2) = 6 (6 × 1) = 6 H (4 × 3) = 12 (6 × 2) = 12

Your Turn! Count the number of atoms of each element on both sides of the arrow to determine whether the following equation is balanced. 2(NH4)3PO4 + 3Ba(C2H3O2)2 → Ba3(PO4)2 + 6NH4C2H3O2 Reactants Products N (2 × 3) = 6 (6 × 1) = 6 H (2×3×4)+(3×3×2) = 42 (6×4) + (6×3) = 42 O (2×4) + (3×2×2) = 20 (2×4) + (6×2) = 20 P (2 × 1) = 2 Ba (3 × 1) = 3

Ions and Ionic Compounds
Transfer of one or more electrons from one atom to another Form electrically charged particles Ionic compound Compound composed of ions Formed from metal and nonmetal Infinite array of alternating Na+ and Cl– ions Formula unit Smallest neutral unit of ionic compound Smallest whole-number ratio of ions 57

Formation of Ionic Compounds
Metal + Non-metal  ionic compound 2Na(s) + Cl2(g)  2NaCl(s) 58

Ionic Compounds Cations Anions Positively charged ions
Formed from metals Atoms lose electrons e.g., Na has 11 e– and 11 p Anions Negatively charged ions Formed from non-metals Atoms gain electrons e.g., Cl has 17 e– and 17 p Na+ has 10 e– and 11 p Cl– has 18 e– and 17 p 59 59

Experimental Evidence for Ions
Electrical conductivity requires charge movement Ionic compounds: Do not conduct electricity in solid state Do conduct electricity in liquid and aqueous states where ions are free to move Molecular compounds: Do not conduct electricity in any state Molecules are comprised of uncharged particles 60

Ions of Representative Elements
Can use periodic table to predict ion charges When we use North American numbering of groups: Cation positive charge = group number 61

Ions of Representative Elements
Noble gases are especially stable Nonmetals Negative (–) charge on anion = number of spaces you have to move to right to get to noble gas Expected charge on O Move two spaces to right O2– What is expected charge on N? Move three spaces to right N3 – N O F Ne 62 62

Rules For Writing Ionic Formulas
Cation given first in formula Subscripts in formula must produce electrically neutral formula unit Subscripts must be smallest whole numbers possible Divide by 2 if all subscripts are even May have to repeat several times Charges on ions not included in finished formula unit of substance If no subscript, then 1 implied 63 63

Determining Ionic Formulas
Example: Formula of ionic compound formed when magnesium reacts with oxygen Mg is group 2A Forms +2 ion or Mg2+ O is group 6A Forms –2 ion or O2– To get electrically neutral particle need 1:1 ratio of Mg2+ and O2– Formula: MgO 64

Determining Ionic Formulas
“Criss-cross” rule Make magnitude of charge on one ion into subscript for other When doing this, make sure that subscripts are reduced to lowest whole number. Example: What is the formula of ionic compound formed between aluminum and oxygen ions? Al3+ O2– Al2O3 65 65

Your Turn! Which of the following is the correct formula for the formula unit composed of potassium and oxygen ions? KO KO2 K2O P2O3 K2O2 66

Your Turn! Which of the following is the correct formula for the formula unit composed of Fe3+ and sulfide ions? FeS Fe3S2 FeS3 Fe2S3 Fe4S6 67

Your Turn! Which of the following is the correct formula for the formula unit composed of ions of magnesium and nitrogen? MgN Mg3N2 Mn3N2 N3Mg2 Mn2N3 68

Cations of Transition Metals
Center (shorter) region of periodic table Much less reactive than group 1A and 2A Still transfer electrons to nonmetals to form ionic compounds number of electrons transferred less clear Form more than one positive ion Can form more than one compound with same non-metal e.g., Fe + Cl FeCl2 and FeCl3 69

Cations of Post-transition Metals
Nine metals Ga, In, Sn, Tl, Pb, Bi, Uut, Uuq, Uub After transition metals and before metalloids Two very important ones – tin (Sn) and lead (Pb) Both have two possible oxidation states Both form two compounds with same nonmetal e.g., Ionic compounds of tin and oxygen are SnO and SnO2 Bismuth Only has +3 charge Bi3+ 70

Ions of Some Transition Metals and Post-transition Metals
71 71

Compounds with Polyatomic Ions
Binary compounds Compounds formed from two different elements Polyatomic ions Ions composed of two or more atoms linked by molecular bonds If ions are negative, they have too many electrons If ions are positive, they have too few electrons Formulas for ionic compounds containing polyatomic ions Follow same rules as ionic compounds Polyatomic ions are expressed in parentheses 72

Polyatomic Ions 73 73

Learning Check Ex. What is the formula of the ionic compound formed between ammonium and phosphate ions? Ammonium = NH4+ Phosphate = PO43– Ex. Between strontium ion and nitrate ion? Strontium = Sr2+ Nitrate = NO32– (NH4)+ (PO4)3– (NH4)3PO4 Sr2+ (NO3)– Sr(NO3)2 74 74

Nomenclature (Naming)
IUPAC system to standardize name of chemical compounds One system so that anyone can reconstruct formula from name We will look at naming ionic compounds of Representative metals Transition metals Monatomic ions Polyatomic ions Hydrates 75 75

Naming Ionic Compounds
Cations: Metal that forms only one positive ion Cation name = English name for metal Na+ sodium Ca2+ calcium Metal that forms more than one positive ion Use Stock System Cation name = English name followed by numerical value of charge written as Roman numeral in parentheses (no spaces) Transition metal Cr2+ chromium(II) Cr3+ chromium(III) 76 76

Naming Ionic Compounds
Anions: Monatomic anions named by adding “–ide” suffix to stem name for element Polyatomic ions use names in Table 2.4 77 77

Learning Check: Name The Following
K2O NH4ClO3 Mg(C2H3O2)2 Cr2O3 ZnBr2 potassium oxide ammonium chlorate magnesium acetate chromium(III) oxide zinc bromide Note: chromium is a transition metal that has a variable charge, hence we must specify. Since Zn has only 1 common charge, this is unnecessary. Other transition metals with only one common charge are Cd2+, and Ag+ 78 78

Learning Check: Determine The Formula
Calcium hydroxide Ca(OH)2 Manganese(II) bromide MnBr2 Ammonium phosphate (NH4)3PO4 Mercury(I) nitride (Hg2)3N2 79 79

Your Turn! Which is the correct name for Cu2S? copper(I) sulfide
copper sulfide copper(II) sulfide copper(II) sulfate copper(I) sulfide copper(I) sulfite 80

Your Turn! Which is the correct formula for ammonium sulfite?
A. NH4SO3 B. (NH4)2SO3 C. (NH4)2SO4 D. NH4S E. (NH4)2S 81

Naming Hydrates Ionic compounds Naming mono- = 1 hexa- 6 di- 2 hepta-
Crystals contain water molecules Fixed proportions relative to ionic substance Naming Name ionic compound Give number of water molecules in formula using Greek prefixes mono- = 1 hexa- 6 di- 2 hepta- 7 tri- 3 octa- 8 tetra- 4 nona- 9 penta- 5 deca- 10 82

Learning Check: Naming Hydrates
CaSO4·2H2O calcium sulfate dihydrate CoCl2·6H2O cobalt(II) chloride hexahydrate FeI3·3H2O iron(III) iodide trihydrate 83

Your Turn! What is the correct formula for copper(II) sulfate pentahydrate? CuSO4 · 6H2O CuSO3 · 5H2O CoSO4 · 4H2O CoSO3 · 5H2O CuSO4 · 5H2O 84

Your Turn! What is the correct name for Fe(NO3)3·9H2O
iron nitrate nonahydrate iron(III) nitrate nonahydrate Ferium (III) nitrate decahydrate iron(III) nitrite nonahydrate iron(III) nitrate heptahydrate 85

Molecular Compounds Molecules Chemical bonds Molecular formulas
Electrically neutral particle Consists of two or more atoms Chemical bonds Attractions that hold atoms together in molecules Arise from sharing electrons between two atoms Group of atoms that make up molecule behave as single particle Molecular formulas Describe composition of molecule Specify number of each type of atom present 86 86

Molecules vs. Ionic Compounds
Discrete unit Water = two hydrogen atoms bonded to one oxygen atom Ionic Compounds Ions packed as close as possible to each other Sodium chloride: Six anions surround each cation; six cations surround each anion No one ion “belongs” to another 87

Molecular Compounds Formed when nonmetals combine
C + O2  CO H2 + O2  2H2O Millions of compounds can form from a few non-metals Organic chemistry and biochemistry Deal with chemistry of carbon + hydrogen, nitrogen, and oxygen A few compounds have only two atoms Diatomics: HCl, CO, HF, NO Most molecules are far more complex Sucrose (C12H22O11) urea (CON2H4) 88

Hydrogen-containing Compounds
Nonmetal hydrides Molecule containing nonmetal + hydrogen Number of hydrogens that combine with nonmetal = number of spaces from nonmetal to noble gas in periodic table N O F Ne 89

3-D Shapes of Molecules Space filling models
Used to give shapes of simple nonmetal hydrides Blue = nitrogen Red = oxygen Yellow = fluorine White = hydrogen 90

Organic Compounds Carbon compounds
Carbon, hydrogen, oxygen, and nitrogen Originally thought these compounds only came from living organisms Now more general Hydrocarbons Simplest organic compounds Contain only C and H 91

Hydrocarbons Belonging to the Alkane Series
92

Alkanes Boiling point increases as number of carbon atoms increases
Space filling models of alkanes Black = carbon White = hydrogen 93

Your Turn! What is the correct name for C4H10? methane ethane propane
D. butane pentane 94

Other Hydrocarbons Alkenes Alkynes
Hydrocarbons with two less H’s than alkanes CnH2n e.g., C2H4 = ethene (ethylene) Replace –ane ending of alkanes with –ene for alkenes Alkynes Hydrocarbons with four fewer H’s than alkanes CnH2n – 2 e.g., C2H2 = ethyne (acetylene) Replace –ane ending of alkanes with –yne for alkynes 95

Other Organic Compounds
Hydrocarbons are basic building blocks of organic chemistry Many other classes of compounds derived from them Alcohols Replace H in alkane with –OH group e.g., CH3OH = methanol (methyl alcohol) C2H5OH = ethanol (ethyl alcohol) Replace –e in an alkane name with –ol Note that ethene is commonly referred to ethylene. 96 96

Your Turn! What is the name of C4H9OH? hexanol propanol pentanol
tetranol butanol 97

Your Turn! What is the name of CH3CH2CH2CH2CCH (i.e., C6H10)? hexene
heptacarbon decahydrogen hexane hexyne heptene 98

Writing Formulas for Organic Compounds
Molecular formula Indicates number of each type of atom in molecule e.g., C2H6 for ethane or C3H8 for propane Order of atoms Carbon Hydrogen Other atoms alphabetically e.g., sucrose is C12H22O11 Emphasize alcohol – write OH group last C2H5OH Structural formula Indicate how carbon atoms are connected Ethane = CH3CH3 Propane = CH3CH2CH3 99

Your Turn! Octane is a hydrocarbon with eight C atoms that is the major component of gasoline. What is the correct molecular formula for octane? C8H14 C8H16 C8H18 C8H17OH C8H15OH 100

Your Turn! What is the correct structural formula for octane? A. CH3CH2CH2CH2CH2CH2CH2CH3 B. CH3CH2CH2CH2CH2CH2CH3 C. C8H18 D. CH3CH2CH2CH2CH2CH2CH2CH2CH3 E. CH3CH2CH2CH2CH2CH2CH2CH2OH 101

Nomenclature of Molecular Compounds
Goal is a name that translates clearly into molecular formula Naming Binary Molecular Compounds Which two elements present? How many of each? Format: First element in formula Use English name Second element Use stem and append suffix –ide Use Greek number prefixes to specify how many atoms of each element 102 102

Naming Binary Molecular Compounds
hydrogen chloride phosphorous pentachloride triselenium dinitride Mono always omitted on first element Often omitted on second element unless more than one combination of same two elements e.g., Carbon monoxide CO Carbon dioxide CO2 When prefix ends in vowel similar to start of element name, drop prefix vowel 1 H 1 Cl HCl 1 P 5Cl PCl5 3 Se 2N Se3N2 103

Learning Check: Name Each
Format: Number prefix + first element name Number prefix + stem + –ide for second element AsF3 = HBr = N2O4 = N2O5 = CO = CO2 = arsenic trifluoride hydrogen bromide dinitrogen tetroxide dinitrogen pentoxide carbon monoxide carbon dioxide 104 104

Your Turn! Which is the correct formula for nitrogen triiodide? N3I
NIO3 N(IO3)3 none of the above 105

Your Turn! Which is the correct name for P4O10? phosphorus oxide
phosphorous decoxide tetraphosphorus decoxide tetraphosphorus oxide decoxygen tetraphosphide 106 106

Your Turn! Which is the correct formula for disulfur decafluoride?
Su2F10 SF5 S2F10 S3F10 S2F4 107

Exceptions to Naming Binary Molecules
Binary compounds of nonmetals + hydrogen No prefixes to be used Get number of hydrogens for each nonmetal from periodic table Hydrogen sulfide = H2S Hydrogen telluride = H2Te Molecules with Common Names Some molecules have names that predate IUPAC systematic names Water H2O ▪ Sucrose C12H22O11 Ammonia NH3 ▪ Phosphine PH3 108

Summary of Naming 109

Chapter 2 Elements, Compounds, and the Periodic Table

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Chapter 2 Elements, Compounds, and the Periodic Table

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