chemical compounds - faculty

CHAPTER 3Chemical Compounds

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ObjectivesYou will be able to do the following.

1. Given a description of a form of matter, classify as an element, compound, or mixture.

2. Write a description of the polar covalent bond between the hydrogen atom and the chlorine atom in HCl. Your description should include a rough sketch of the electron cloud that represents the electrons involved in the bond.

3. Write a description of the process that leads to the formation of the ionic bond between sodium and chlorine atoms in sodium chloride.

4. Write a description of the sodium chloride crystal structure.5. Write a description of the difference between a nonpolar covalent bond, a polar

covalent bond, and an ionic bond. Your description should include rough sketches of the electron clouds that represent the electrons involved in the formation of the each bond.

6. Given the names or formulas for two elements, identify the bond that would form between them as covalent or ionic.

7. Given a table of electronegativities, do the following. (See pages 354-357 of the text.)

a. Classify chemical bonds as nonpolar covalent, polar covalent, or ionic. b. Identify which of two atoms in a polar covalent bond has a partial negative

charge and which atom has a partial positive charge. c. Identify which of two atoms in an ionic bond has a negative charge and

which atom has a positive charge.d. Given two bonds, determine which of the bonds would be expected to be

more polar. 8. Identify the most common number of covalent bonds that each of the following

elements form: hydrogen, Group 17 (halogens), oxygen, nitrogen, and carbon.9. Convert between the description of the number of atoms of each element found

in a compound and its chemical formula. 10. Draw Lewis structures from chemical formulas for compounds that have all of

their atoms with their most common bonding pattern.11. Describe the molecular geometry of CH4, NH3, and H2O molecules.12. Write a description of the attractions between water molecules in liquid and solid

water. 13. Write a description of the structure of liquid water. Your description should

include a sketch of the particles in the liquid. 14. Convert between the names and chemical formulas for water, ammonia,

methane, ethane, propane, butane, pentane, and hexane. 15. Write or identify prefixes for the numbers 1-10. (For example, mono- represents

one, di- represents two, etc.)

32 Chapter 3 Chemical Compounds

16. Write or identify the roots of the nonmetals. (For example, the root for oxygen is ox−.)

17. Convert between the systematic names and chemical formulas for binary covalent compounds.

18. Convert between the complete name, the common name, and the chemical formula for HF, HCl, HBr, HI, and H2S.

19. Write or identify the ionic charges that form for the following elements: group 17 (halogens), oxygen, sulfur, selenium, nitrogen, phosphorus, hydrogen, group 1 (alkali metals), group 2 (alkaline earth metals), group 3 elements, aluminum, iron, silver, copper, and zinc.

20. Write an explanation for why the elements listed in the previous objective (except for (iron, silver, copper, and zinc) form their ionic charges.

21. Convert between the names and chemical formulas for the monatomic ions. 22. Given the common name, write the chemical formula for the following ions:

ferric, ferrous, cupric, and cuprous.23. Convert between the names and chemical formulas for the common polyatomic

ions hydroxide, ammonium, cyanide, acetate, dichromate, permanganate, oxalate, sulfate, nitrate, phosphate, carbonate, chlorate, chromate, bromate, and iodate.

24. Convert between the names and chemical formulas for the per(root)ate, (root)ite, and hypo(root)ite oxyanions that have the same element as the (root)ate ions.

25. Convert between the names and chemical formulas for the polyatomic anions that are derived from the additions of H+ ions to anions with −2 or −3 charges. For example, H2PO4

− is dihydrogen phosphate. 26. Given the common name bicarbonate, bisulfate, or bisulfite, write its chemical

formula.27. Convert between the names and chemical formulas for ionic compounds.28. Convert between the name or formula of a (root)ate anion and the name or

formula of the (root)ic acid that forms from the (root)ate anion. 29. Write or identify a description of alcohols. 30. Convert between the correct name and chemical formula for each of the

following alcohols: methanol, ethanol, and 2-propanol. 31. Given one of the following common names for alcohols, write its chemical

formula: methyl alcohol, ethyl alcohol, and isopropyl alcohol.32. Convert between the name and chemical formula the sugars glucose and sucrose. 33. Given a name or chemical formula, identify whether it represents a binary

ionic compound, an ionic compound with polyatomic ion(s), a binary covalent compound, a binary acid, an oxyacid, a common alcohol, or a common sugar.

34. Convert between the name and chemical formula for binary ionic compounds, ionic compounds with polyatomic ion(s), binary covalent compounds, binary acids, oxyacids, common alcohols, and common sugars.

35. Given a description of a compound, identify it as either a molecular compound or an ionic compound.

For objectives 36-42, see pages 409-416 of the text.36. Write or identify a description of dipole-dipole attractions between polar

molecules. 37. Draw a sketch of hydrogen chloride molecules in the liquid form showing the

dipole-dipole attractions that hold the particles together.

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38. Write or identify a description of hydrogen bonds between molecules of HF, H2O, and NH3.

39. Write or identify a description of London dispersion forces between nonpolar molecules. Your description should include mention of how the attractions form beginning with nonpolar molecules.

40. Write a description of how London forces form between polar molecules.41. Write an explanation for why larger molecules have stronger London forces. 42. Write or identify the nature of the particles and the nature of the attractions

between the particles in the liquid and solid form for metallic elements, the noble gases, carbon in the diamond form, the other nonmetallic elements (hydrogen, nitrogen, oxygen, sulfur, selenium, phosphorus, and the halogens), any ionic compound, any hydrocarbon, hydrogen chloride, hydrogen fluoride, water, alcohols, and ammonia.

43. Draw constitutional isomers of C4H10, C5H12, and C6H14. 44. Convert between names and formulas for alkanes, including branched alkanes

with alkyl groups in Table 3.5 of the text. 45. Write or identify the two abbreviations associated with atomic (unified) mass

unit. 46. Given a periodic table that shows atomic masses of the elements, convert between

mass of an element and moles of that element. 47. Given a formula for a molecular substance and a periodic table that includes

atomic masses for the elements, calculate the substance’s molecular mass.48. Given enough information to calculate a molecular substance’s molecular mass,

convert between mass and moles of the substance. 49. Write a description of the similarities and differences between the term molecule

and the term formula unit. 50. Given a formula for an ionic compound and a periodic table that includes atomic

masses for the elements, calculate the compound’s formula mass.51. Given enough information to calculate an ionic compound’s formula mass,

convert between mass and moles of the compound. 52. Given a chemical formula for a compound, write conversion factors that convert

between moles of compound and moles of element in the compound. 53. Make conversion between mass of compound and mass of element in the

compound. 54. Given a chemical formula for a compound, calculate the percentage of the

elements in the compound. 55. Convert between the definition and the term for the following words or phrases.

Skip sections 3.10 and 3.12 in the text.


Chapter 3 GlossaryElement A substance that cannot be chemically converted into simpler substances; a

substance in which all of the atoms have the same number of protons and therefore the same chemical characteristics.

Compound A substance that contains two or more elements, the atoms of these elements always combining in the same whole-number ratio.

Pure substance A sample of matter that has constant composition. There are two types of pure substances, elements and compounds.

Mixture A sample of matter that contains two or more pure substances and has variable composition.

Chemical bond An attraction between atoms or ions in chemical compounds. Covalent bonds and ionic bonds are examples.

Polar covalent bond A covalent bond in which electrons are shared unequally, leading to a partial negative charge on the atom that attracts the electrons more and to a partial positive charge on the other atom.

Nonpolar covalent bond A covalent bond in which the difference in electron-attracting ability of two atoms in a bond is negligible (or zero), so the atoms in the bond have no significant charges.

Ion Any charged particle, whether positively or negatively charged. Cation An ion formed from an atom that has lost one or more electrons and thus has

become positively charged. Anion An ion formed from an atom that has gained one or more electrons and thus

has become negatively charged. Ionic bond The attraction between a cation and an anion. Ionic hydrate Ionic compounds with water molecules trapped within the crystal

lattice. Water of hydration The associated water in ionic hydrates. Electronegativity A measure of the electron-attracting ability of an atom in a chemical

bond. Molecular compound A compound composed of molecules. In such compounds, all

of the bonds between atoms are covalent bonds. Ionic compound A compound that consists of ions held together by ionic bonds. Chemical formula A concise written description of the components of a chemical

compound. It identifies the elements in the compound by their symbols and indicates the relative number of atoms of each element with subscripts.

Empirical formula A chemical formula that includes positive integers that describe the simplest ratio of the atoms of each element in a compound.

Molecular formula The chemical formula that describes the actual numbers of atoms of each element in a molecule of a compound.

Valence electrons The electrons that are most important in the formation of chemical bonds. The number of valence electrons for the atoms of an element is equal to the element’s A-group number on the periodic table.

Electron-dot symbol A representation of an atom that consists of its elemental symbol surrounded by dots representing its valence electrons.

Lewis structure A representation of a molecule that consists of the elemental symbol for each atom in the molecule, lines to show covalent bonds, and pairs of dots to indicate lone pairs.

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Double bond A link between atoms that results from the sharing of four electrons. It can be viewed as two 2-electron covalent bonds.

Triple bond A link between atoms that results from the sharing of six electrons. It can be viewed as three 2-electron covalent bonds.

Lone pair Two electrons that are not involved in the covalent bonds between atoms but are important for explaining the arrangement of atoms in molecules. They are represented by pairs of dots in Lewis structures.

Bond angle The angle formed by straight lines (representing bonds) connecting the nuclei of three adjacent atoms.

Tetrahedral The molecular shape that keeps the negative charge of four electron groups as far apart as possible. This shape has angles of 109.5° between the atoms.

Binary covalent compound A compound composed of two nonmetallic elements.Organic chemistry The branch of chemistry that involves the study of carbon-based

compounds. Organic compound A carbon-based compound. Hydrocarbons Compounds that contain only carbon and hydrogen. Alcohols Compounds that contain a hydrocarbon group with one or more -OH

groups attached. Monatomic anions Negatively charged particles, such as Cl−, O2−, and N3−, that

contain single atoms with a negative charge. Monatomic cations Positively charged particles, such as Na+, Ca2+, and Al3+, that

contain single atoms with a positive charge. Binary ionic compound An ionic compound whose formula contains one symbol

for a metal and one symbol for a nonmetal.Polyatomic ion A charged collection of atoms held together by covalent bonds. Oxyanion A polyatomic ions with the general formula HaXbOc

d−. (The a can be 0.)Binary acid Substances that have the general formula of HX(aq), where X is one of

the first four halogens HF(aq), HCl(aq), HBr(aq), and HI(aq).Oxyacids (or oxoacids) Molecular substances that have the general formula HaXbOc.

In other words, they contain hydrogen, oxygen, and one other element represented by X; the a, b, and c represent subscripts.

Intermolecular attraction Attraction between molecules. Dipole A molecule that contains an asymmetrical distribution of positive and negative

charges. Dipole-dipole attraction The intermolecular attraction between the partial negative

end of one polar molecule and the partial positive end of another polar molecule. Hydrogen bond The intermolecular attraction between a nitrogen, oxygen, or

fluorine atom of one molecule and a hydrogen atom bonded to a nitrogen, oxygen, or fluorine atom in another molecule.

Metallic bond The attraction between the positive metal cations that form the basic structure of a solid metal and the negative charge from the mobile sea of electrons that surround the cations.

London dispersion forces, London forces, or dispersion forces The attractions produced between molecules by instantaneous and induced dipoles.

Isomers Compounds that have the same molecular formula but different molecular structures.


Constitutional isomers (also called structural isomers) Compounds with the same molecular formula that differ in the order in which their atoms are bonded together.

Alkanes Hydrocarbons (compounds composed of carbon and hydrogen) in which all of the carbon-carbon bonds are single bonds.

Weighted average mass A mass calculated by multiplying the decimal fraction of each component in a sample by its mass and adding the results of each multiplication together.

Atomic mass unit 1/12 the mass of a carbon-12 atom. It is sometimes called a unified mass unit. Its accepted abbreviation is u, but amu is sometimes used.

Atomic mass The weighted average of the masses of the naturally occurring isotopes of an element.

Mole The amount of substance that contains the same number of particles as there are atoms in 12 g of carbon-12.

Avogadro’s number The number of atoms in 12 g of carbon-12. To four significant figures, it is 6.022 x 1023.

Molar mass The mass in grams of one mole of substance. (The number of grams in the molar mass of an element is the same as its atomic mass. The number of grams in the molar mass of a molecular compound is the same as its molecular mass. The number of grams in the molar mass of an ionic compound is the same as its formula mass.)

Molecular mass The weighted average of the masses of the naturally occurring molecules of a molecular substance. It is the sum of the atomic masses of the atoms in a molecule.

Formula unit A group represented by a substance’s chemical formula, that is, a group containing the kinds and numbers of atoms or ions listed in the chemical formula. It is a general term that can be used in reference to elements, molecular compounds, or ionic compounds.

Formula mass The weighted average of the masses of the naturally occurring formula units of the substance. It is the sum of the atomic masses of the atoms in a formula unit.

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Figure 3.2Ionic Bond Formation

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Figure 3.1Elements Versus Compounds


Summary of Covalent and Ionic Bond Formation When atoms of different elements form chemical bonds, the electrons in the bonds

can shift from one bonding atom to another. The atom that attracts electrons more strongly will acquire a negative charge, and

the other atom will acquire a positive charge. The more the atoms differ in their electron-attracting ability, the more the electron

cloud shifts from one atom toward another. If there is a large enough difference in electron-attracting ability, 1, 2, or 3 electrons

can be viewed as shifting completely from one atom to another. The atoms become positive and negative ions, and the attraction between them is called an ionic bond.

If the electron transfer is significant but not enough to form ions, the atoms acquire partial positive and partial negative charges. The bond in this situation is called a polar covalent bond.

If there is no shift of electrons or if the shift is negligible, no significant charges will form, and the bond will be a nonpolar covalent bond.

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Figure 3.3 Classifying Compounds

EXERCISE 3.1 - Classifying Compounds

Classify each of the following substances as either a molecular compound or an ionic compound.

a. formaldehyde, CH2O (used in embalming fluids)b. magnesium chloride, MgCl2 (used in fireproofing wood and in paper

manufacturing)

Table 3.1 Guidelines for Classifying Bonds as Covalent or Ionic

Combination Type of bond Examples

Metal-Nonmetal Usually Ionic NaCl or MgO

Nonmetal-Nonmetal Covalent HCl or CO

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2

3

4

5

6

7

11 2

3 4 5 6 7 8 9 10 11 12

13 14 15 16 17

18

1A

8A

7A6A5A4A3A2A

3B 2B1B8B8B8B7B6B5B4B

2.04 2.55 3.04 3.44 3.98

1.61 1.90 2.19 2.58 3.16

1.81 2.01 2.18 2.55 2.96

1.78 1.96 2.05 2.1 2.66

B C N O F

Al Si P S Cl

Ga Ge As Se Br

In Sn Sb Te I1.62 2.33 2.02 2.0Tl Pb Bi Po

2.20

1.65

1.69

H

Zn

Cd2.00Hg

1.88 1.91 1.90

2.28 2.20 1.93Co Ni Cu

Rh Pd Ag2.20 2.28 2.54Ir Pt Au

1.66 1.55 1.83

2.16 1.9 2.2Cr Mn Fe

Mo Tc Ru2.36 1.9 2.2W Re Os

1.36 1.54 1.63

1.22 1.33 1.6Sc Ti V

Y Zr Nb1.27 1.3 1.5Lu Hf Ta

0.98 1.57

0.93 1.31

0.82 1.00

0.82 0.95

Li Be

Na Mg

K Ca

Rb Sr0.79 0.89

0.7 0.9

Cs Ba

Fr Ra

2.6Xe

2.2At

3.00Kr

Figure 3.4Electronegativities

Sample Study Sheet 3.1: Electronegativity, Types of Chemical Bonds, and Bond Polarity

TIP-OFF You wish to (1) classify a chemical bond as nonpolar covalent, polar covalent, or ionic, (2) identify which element in a polar covalent bond is partially negative and which is partially positive, (3) identify which element in an ionic bond forms the anion and which forms the cation, or (4) identify which of two bonds is more polar.

GENERAL STEPS

• Use the following guidelines to identify a chemical bond as ionic, nonpolar covalent, or polar covalent. (If both atoms are nonmetals, the bond is covalent.)

∆EN < 0.4 → Nonpolar Covalent

∆EN 0.4-1.7 → Polar Covalent

∆EN > 1.7 → Ionic

• Use the following guidelines to identify which element in a polar covalent bond is partially negative and which is partially positive.

Higher electronegativity → partial negative charge

Lower electronegativity → partial positive charge

• Use the following guidelines to identify which element in an ionic bond forms the anion and which forms the cation.

Nonmetal, which has a higher electronegativity → anion

Metal, which has the lower electronegativity → cation

• Use the following guideline to decide which of two bonds is more polar.

The greater the ∆EN is, the more polar the bond.


Figure 3.5Electronegativities and Bond Type

EXERCISE 3.2 - Electronegativities and Bond Type

Classify the following bonds as nonpolar covalent, polar covalent, or ionic. If a bond is polar covalent, identify which atom has the partial negative charge and which has the partial positive charge. If a bond is ionic, identify which atom has the negative charge and which has the positive charge.

a. N bonded to H

b. N bonded to Cl

c. Ca bonded to O

d. P bonded to F

EXERCISE 3.3 - Electronegativities and Bond Polarity

Which bond would you expect to be more polar, P-H or P-F?

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Table 3.2Electron-Dot Symbols and Usual Numbers of Bonds and Lone Pairs for Nonmetallic Elements

Group 4A Group 5A Group 6A Group 7A4 valence electrons

X

5 valence electrons

X

6 valence electrons

X

7 valence electrons

X

4 bonds No lone pairs

3 bonds 1 lone pair

2 bonds 2 lone pairs

1 bond 3 lone pairs

carbon-C

C

nitrogen-N

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EXERCISE 3.4 - Drawing Lewis Structures from Formulas

Draw a Lewis structure for each of the following formulas:

a. nitrogen triiodide, NI3 (explodes at the slightest touch)

b. hexachloroethane, C2Cl6 (used to make explosives)

c. hydrogen peroxide, H2O2 (a common antiseptic)

d. ethylene (or ethene), C2H4 (used to make polyethylene)


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Figure 3.6 Three Ways to Describe a Methane Molecule

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Figure 3.7 Three Ways to Describe an Ammonia Molecule

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Figure 3.8 Three ways to Describe a Water Molecule

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Figure 3.9Attractions Between water Molecules

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Figure 3.10Liquid Water

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A Shockwave animation on our website will help you to visualize the structure

of water:

http://www.mpcfaculty.net/mark_bishop/water.htm

Table 3.3 Names to Memorize for Some Binary Covalent Compounds

Name Formula Name Formula

water H2O ammonia NH3

methane CH4 ethane C2H6

propane C3H8 butane C4H10

pentane C5H12 hexane C6H14

Follow these steps to write the names for binary covalent compounds. • If the subscript for the first element is greater than one, indicate the identity

of the subscript using prefixes from the table below. We do not write mono- at the beginning of a compound’s name.

Example: We start the name for N2O3 with di-.• Attach the selected prefix to the name of the first element in the formula. If

no prefix is to be used, begin with the name of the first element.Example: We indicate the N2 portion of N2O3 with dinitrogen.

• Select a prefix to identify the subscript for the second element (even if its subscript is understood to be one). Leave the “a” off the end of the prefixes that end in “a” and the “o” off of mono- if they are placed in front of an element whose name begins with a vowel (oxygen or iodine).

Example: The name of N2O3 grows to dinitrogen tri-.

• Write the root of the name of the second element in the formula. Example: The name of N2O3 becomes dinitrogen triox-.

• Add -ide to the end of the name.Example: The name of N2O3 is dinitrogen trioxide.

Table 3.4 Prefixes Used in the Names of Binary Covalent Compounds to Indicate the Number of Atoms of each Element in the Formula

Number of atoms

Prefix Number of atoms

Prefix

1 mon(o)- 6 hex(a)-2 di- 7 hept(a)-3 tri- 8 oct(a)-4 tetr(a)- 9 non(a)-5 pent(a)- 10 dec(a)-

http://www.mpcfaculty.net/mark_bishop/water.htm


Table 3.5 Roots for the Nonmetals

Element Root Element Root Element Root Element RootC carb- N nitr- O ox- F fluor-

P phosph- S sulf- Cl chlor-As arsen- Se selen- Br brom-

I iod-

Table 3.6 Names for Binary Covalent Compounds with Atoms that Combine in Only

One Ratio

Formula Complete Name Common NameHF hydrogen monofluoride hydrogen fluorideHCl hydrogen monochloride hydrogen chlorideHBr hydrogen monobromide hydrogen bromideHI hydrogen moniodide hydrogen iodide

H2S dihydrogen monosulfide (or dihydrogen sulfide)

hydrogen sulfide

The first step in writing formulas when given the systematic name of a binary covalent compound is to recognize the name as representing a binary covalent compound. It will have one of the following general forms.

prefix(name of nonmetal) prefix(root of name of nonmetal)ide (e.g. dinitrogen pentoxide)or (name of nonmetal) prefix(root of name of nonmetal)ide (e.g. carbon dioxide)or (name of nonmetal) (root of nonmetal)ide (e.g. hydrogen fluoride)

Follow these steps for writing formulas for binary covalent compounds when you are given a systematic name. Notice that they are the reverse of the steps for writing names from chemical formulas.

• Write the symbols for the elements in the order mentioned in the name. • Write subscripts indicated by the prefixes. If the first part of the name has no

prefix, assume it is mono-.

EXERCISE 3.5 - Naming of Binary Covalent Compounds

Write names that correspond to the following formulas: (a) P2O5, (b) PCl3, (c) CO, (d) H2S, and (e) NH3.

EXERCISE 3.6 - Writing Formulas for Binary Covalent Compounds

Write formulas that correspond to the following names: (a) disulfur decafluoride, (b) nitrogen trifluoride, (c) propane, and (d) hydrogen chloride.

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Figure 3.11Common Monatomic Ions

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Anion Name Anion Name Anion NameN3− nitride O2− oxide H− hydrideP3− phosphide S2− sulfide F− fluoride

Se2− selenide Cl− chlorideBr− bromideI− iodide

Table 3.7Names of the Monatomic Anions

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Figure 3.12 Monatomic Anion Charges


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Figure 3.13 Monatomic Anion Charges

EXERCISE 3.7 - Naming Monatomic Ions

Write names that correspond to the following formulas for monatomic ions: (a) Mg2+,

(b) F−, and (c) Sn2+.

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EXERCISE 3.8 - Formulas for Monatomic Ions

Write formulas that correspond to the following names for monatomic ions: (a) bromide ion, (b) aluminum ion, and (c) gold(I) ion.

Figure 3.14 Sodium Chloride Structure

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Table 3.8 Ionic Charges to Memorize

Element Charge Element Charge

iron 2+ or 3+ copper 1+ or 2+zinc 2+ cadmium 2+silver 1+ (and very rarely +2)

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Figure 3.15Cesium Chloride Crystal Structure

Table 3.9 Naming Metals with Two Possible Charges

Ion Systematic (Preferred) Name

Nonsystematic Name

Fe2+ iron(II) ferrous

Fe3+ iron(III) ferric

Cu+ copper(I) cuprous

Cu2+ copper(II) cupric

Table 3.10 Polyatomic ions with common names.

Ion Name Ion Name

OH− hydroxide NH4+ ammonium

CN− cyanide C2H3O2− acetate

C2O42− oxalate MnO4

− permanganate

Cr2O72− dichromate


Table 3.11 Convention for Naming Oxyanions

Relationship General Name Example Name Example Formula

one more oxygen atom than (root)ate

per(root)ate perchlorate ClO4−

(root)ate chlorate ClO3−

one less oxygen atom than (root)ate

(root)ite chlorite ClO2−

two less oxygen atoms than (root)ate

hypo(root)ite hypochlorite ClO−

Table 3.12 (Root)ate Polyatomic Ions

Ion Name Ion Name

SO42− sulfate NO3

− nitrate

PO43− phosphate CO3

2− carbonate

ClO3− chlorate CrO4

2− chromate

BrO3− bromate IO3

− iodate

Table 3.13 Formula and Names for Polyatomic Ions Containing Hydrogen

Formula Name Formula Name

HCO3− hydrogen carbonate HPO4

2− hydrogen phosphate

HS− hydrogen sulfide H2PO4− dihydrogen phosphate

Table 3.14 Systematic and Nonsystematic Names for Some Polyatomic Ions

Formula Systematic (Preferred) Name Nonsystematic Name

HCO3− hydrogen carbonate bicarbonate

HSO4− hydrogen sulfate bisulfate

HSO3− hydrogen sulfite bisulfite

Note: You should use the systematic name, but it will also be useful to know the nonsystematic names.

Table 3.15 Summary of the Ways that Cations are Named

Type of cation General Name Example

For metals with one possible charge (Groups 1, 2, 3 - Al, Zn, and Cd)

name of metal Mg2+ is magnesium.

For metals with more than one possible charge (The rest of the metals)

name of metal(Roman numeral)

Cu2+ is copper(II).

Polyatomic cations Ammonium is our only example in this category.

NH4+ is

ammonium.

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Table 3.16 Summary of the Ways that Anions are Named

Type of anion General Name Example

For monatomic anions (root of nonmetal)ide O2− is oxide.

For polyatomic anions name of polyatomic ion NO3− is nitrate.

EXERCISE 3.9 - Ionic Formulas to Names

Write the names for the following formulas: LiCl, CaSO4, MnF3, NH4F, Cr2S3,

Mg3(PO3)2, ZnCrO4, and AgBrO2

EXERCISE 3.10 - Ionic Names to Formulas

Write the formulas for the following names: aluminum oxide, ammonium chloride, cobalt(II) sulfide, ferrous sulfate, silver chlorite, ammonium hydrogen phosphate, and calcium bicarbonate.

Table 3.17 Two ways to name binary covalent hydrides

Formula Named as Binary Covalent Formula Named as Binary acid

HF or HF(g)

hydrogen monofluoride or hydrogen fluoride

HF(aq) hydrofluoric acid

HCl or HCl(g)

hydrogen monochloride or hydrogen chloride

HCl(aq) hydrochloric acid

HBr or HBr(g)

hydrogen monobromide or hydrogen bromide

HBr(aq) hydrobromic acid

HI or HI(g)

hydrogen moniodide or hydrogen iodide

HI(aq) hydroiodic acid or hydriodic acid

H2S or H2S(g)

dihydrogen sulfide or hydrogen sulfide

H2S(aq) hydrosulfuric acid


Table 3.18 Relationship Between the (Root)ate Polyatomic Ions and the (Root)ic Acids

Oxyanion formula

Oxyanion name

Oxyacid formula

Oxyacid name

NO3− nitrate HNO3 nitric acid

C2H3O2− acetate HC2H3O2 acetic acid

SO42− sulfate H2SO4 sulfuric acid (Note that the whole name

sulfur is used in the oxyacid name.)CO3

2− carbonate H2CO3 carbonic acidPO4

3− phosphate H3PO4 phosphoric acid (Note that an the root of phosphorus in an oxyacid name is phosphor-.)

ClO3− chlorate HClO3 chloric acid

BrO3− bromate HBrO3 bromic acid

IO3− iodate HIO3 iodic acid

C2O42− oxalate H2C2O4 oxalic acid

CrO42− chromate H2CrO4 chromic acid

Note: When enough H+ ions are added to the (root)ate polyatomic ion to completely neutralize the

charge, the (root)ic acid forms.

Table 3.19 Convention for Naming Oxyacids

Relationship General Name Example Name Example Formula

one more oxygen atom than (root)ic

per(root)ic acid perchloric acid HClO4

(root)ic acid chloric acid HClO3

one less oxygen atom than (root)ic

(root)ous acid chlorous acid HClO2

two less oxygen atoms than (root)ic

hypo(root)ous acid hypochlorous acid HClO

Table 3.20 Names of Common Alcohols (You should memorize them.)

Preferred Name Common Name Formulamethanol methyl alcohol CH3OHethanol ethyl alcohol C2H5OH2-propanol isopropyl alcohol C3H7OH

Table 3.21 Names and Formulas for Common Sugars (You should memorize them.)

Name Formulaglucose C6H12O6

sucrose C12H22O11

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Table 3.22 Identifying Types of Compounds from Formulas

Type of Compound General Formula Example

binary ionic MaAb NaCl

ionic with an oxyanion

MaHbXcOd or (NH4)aHbXcOd (b can be 0)

Li2HPO4 or NaNO3 or NH4NO3

binary covalent AaBb N2O5

binary acid HX(aq) or H2S(aq) HCl(aq) or H2S(aq)

oxyacid HaXbOc H2SO4

hydrocarbon CaHb C2H6

alcohol memorize examples CH3OH, C2H5OH, C3H7OH

sugar memorize examples C6H12O6, C12H22O11

M = symbol of metalA and B = symbols of nonmetalsX = some element other than H or Oa, b, c & d indicate subscripts

Table 3.23 Identifying Types of Compounds from Names

Type of Compound Tip-off Examples

binary ionic -ide ending with metal-nonmetal

sodium chloride

ionic with an oxyanion -ite or -ate ending sodium nitrate

binary covalent

a. with common names recognize name water or methane

b. hydrocarbons There are other possibilities, but those we will see will include: (root)ane

n-pentane

c. Other -ide ending with nonmetal-nonmetal

nitrogen trifluoride

binary acid hydro(root)ic acid hydrochloric acid

oxyacid

per(root)ic acid

(root)ic acid

(root)ous acid

hypo(root)ous acid

perchloric acid

chloric acid

chlorous acid

hypochlorous acid

alcohol name ends in -anol or alcohol

methanol or methyl alcohol

sugar name ends in -ose glucose or sucrose


EXERCISE 3.11 - Nomenclature, Formulas to Names

Write names for the following formulas: P2O5, PCl3, CO, H2S(g), H2S(aq), NH3, H3PO4, H3PO3, H3PO2, CH3OH, and C12H22O11.

EXERCISE 3.12 - Nomenclature, Names to Formulas

Write formulas for the following names: disulfur decafluoride, nitrogen trifluoride, butane, hydrogen chloride, hydrochloric acid, carbonic acid, periodic acid, ethanol or ethyl alcohol, and glucose.

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Figure 3.16London Forces

Figure 3.17Larger Molecules, Stronger London Forces

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Table 3.24 Solid and Liquid Form

Type of Substance Particles to Visualize

Examples Type of Attraction Between Particles in

Solid or LiquidElements Metal Cations in a

sea of electronsAu Metallic Bond

Noble Gases Atoms Xe London ForcesCarbon (diamond) Carbon Atoms C(dia) Covalent BondsOther Nonmetal Elements

Molecules H2, N2, O2, F2, Cl2, Br2, I2, S8, Se8, P4

London Forces

Ionic Compounds

Cations and Anions

NaCl Ionic Bond

Molecular Compounds

Nonpolar Molecular

Molecules Hydrocarbons London Forces

Polar molecules without H-F, O-H, or N-H bond

Molecules HCl Dipole-Dipole Forces

Molecules with H-F, O-H or N-H bond

Molecules HF, H2O, alcohols, NH3

Hydrogen Bonds

Figure 3.19Polar Molecules and London forces

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EXERCISE 3.13 - Types of Particles and Attractions

Complete the following table by (1) writing the name for the type of particle viewed as forming the structure of a solid, liquid, or gas of each of the following substances and (2) writing the name of the type of attraction holding these particles in the solid and liquid form.

Substance Particles to Visualize Type of Attraction

iron

MgO

iodine

CH3OH

NH3

hydrogen chloride

C (diamond)

lithium sulfate

Figure 3.19Types of Attractions

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SymmetricalAsymmetrical

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TIP-OFF - You are converting between moles and mass of element.

GENERAL STEPS - The procedure involves unit analysis using the molar mass as a conversion factor. (See Figure 3.1.)

• The molar mass comes from the atomic mass for an element.

Atomic mass = ∑ (decimal fraction of isotope) (mass of isotope)for each

isotope

• Use the following general form for the conversion factor that comes from molar mass.

Atomic Mass

Molar massMass of element moles of atoms

(atomic mass) g element

1 mol element

Sample Study Sheet 3.3: Calculation of Atomic Mass

Sample Study Sheet 3.4: Molar Mass Calculations for Elements

TIP-OFF - You are asked to calculate atomic mass, and you are given the percentage abundance and mass of each isotope. (The tip-offs are more subtle for other types of calculations.)

GENERAL STEPS

• Write the memorized atomic mass equation with the decimal fraction and mass of each isotope given in the problem. (The decimal fraction can be calculated by dividing the percentage abundance by 100.)

Atomic mass = ∑ (decimal fraction of isotope) (mass of isotope)for each

isotope

Decimal fraction of isotope =percent abundance

100

• Complete the calculation and report your answer.

If you are told the number of decimal positions to report, round your answer accordingly.

If you are not given the number of decimal positions to report, use the rules for significant figures to decide how to round off your answer.

Atomic masses can be reported with the unit u for unified mass unit (or atomic mass unit), but they are often reported without a unit.

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EXERCISE 3.14 - Element Molar Mass Calculation

An analysis of the element lithium shows that 7.5% of the lithium atoms are lithium-6 atoms, and 92.5% are lithium-7 atoms. Each atom of lithium-6 has a mass of 6.0151214 u, and each atom of lithium-7 has a mass of 7.0160030 u.

a. What is the atomic mass of lithium? (Report your answer to the third decimal position, ±0.001.)

b. Write a conversion factor that will convert between grams of the element lithium and moles of lithium.

c. How many moles of lithium are in a sample of lithium that has a mass of 7.249 pounds?

EXERCISE 3.15 - Element Molar Mass Calculation

Gold is often sold in units of troy ounces. (To four significant figures, there are 31.10 grams per troy ounce.) How many moles of gold, Au, are there is 1.00 troy ounce of pure gold?


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Figure 3.20Formula Units

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TIP-OFF - You are converting between moles and mass for a molecular compound.

GENERAL STEPS - The procedure involves unit analysis using the molar mass as a conversion factor.

• Form molecular compounds, the molar mass comes from the molecular mass of a molecular compound.

Molecular mass = ∑ atomic massfor each atom

in a molecule


Sample Study Sheet 3.5: Molar Mass Calculations for Molecular Compounds

Molecular Mass

Molar massMass of molecular compound moles of molecules

(molecular mass) g molecular compound

1 mol molecular compound

EXERCISE 3.16 - Molar Mass and Molecular Compounds

A typical 6.0 fluid ounce glass of wine contains about 16 g of ethanol, C2H5OH. a. What is the molecular mass of C2H5OH?

b. Write a conversion factor that will convert between mass and moles of C2H5OH.

c. What is the volume in milliliters of 1.0 mole of pure C2H5OH? (The density of ethanol is 0.7893 g/mL.)


TIP-OFF - You are converting between moles and mass for an ionic compound.

GENERAL STEPS - The procedure involves unit analysis using the molar mass as a conversion factor.

• The molar mass comes from the atomic mass for an element.

Formula mass = ∑ atomic massfor each atom

in a formulaunit


Sample Study Sheet 3.6 Molar Mass Calculations for Ionic Compounds

EXERCISE 3.17 - Molar Mass and Ionic Compounds

A quarter teaspoon of a typical baking powder contains about 0.4 g of sodium hydrogen carbonate, NaHCO3, which is often called bicarbonate of soda.

a. Calculate the formula mass of sodium hydrogen carbonate.

b. Write a conversion factor that could be used to convert between mass and moles of NaHCO3.

c. How many moles of NaHCO3 are there is 0.4 g of NaHCO3?

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TIP-OFF - When you analyze the type of unit you have and the type of unit you want, you recognize that you are converting between a unit associated with an element and a unit associated with a compound containing that element. GENERAL STEPS

• Convert the given unit to moles of the first substance.

This step often requires converting the given unit into grams, after which the grams can be converted into moles using the molar mass of the substance.

• Convert moles of the first substance to moles of the second substance using the molar ratio derived from the formula for the compound.

You either convert from moles of element to moles of compound or moles of compound to moles of element.

• Convert moles of the second substance to the desired units of the second substance.

This step requires converting moles of the second substance into grams of the second substance using the molar mass of the second substance, after which the grams can be converted to the specific units that you want.

The following describes a shortcut for these problems. Use the following general conversion factor in your unit analysis set-up. Like all conversion factors, this conversion factor can be used in the form

described below or in the inverted form.

Sample Study Sheet 3.7: Converting Between Mass of Element and Mass of Compound Containing the Element

EXERCISE 3.18 - Conversion Between Mass of Element and Mass of Compound

Disulfur dichloride, S2Cl2, is used in vulcanizing rubber and hardening soft woods. It can be made from the reaction of pure sulfur with chlorine gas. What is the mass of S2Cl2 that contains 1.238 kg S?

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Figure 3.21General Steps for Converting between the Mass of an Element and the Mass of a Compound Containing the Element The calculation can be set up to convert from the mass of an element to the mass of a compound (top to bottom) or from the mass of a compound to the mass of an element (bottom to top).

EXERCISE 3.19 - Conversion Between Mass of element and Mass of Compound

Vanadium metal, used as a catalyst and to make steel, is produced from the reaction of vanadium(V) oxide, V2O5, and calcium metal. What is the mass in kilograms of vanadium in 2.3 metric tons of V2O5?

EXERCISE 3.20 - Conversion Between Mass of element and Mass of Compound

Calamine has two definitions. It is a naturally occurring zinc silicate that has the equivalent of 67.5% zinc oxide, ZnO, and it is a substance that is used to make the calamine lotion. The calamine used for the lotion is 98% ZnO. The naturally occurring calamine is used to make zinc metal. What is the maximum mass in kilograms of zinc in 1.347 x 104 kg of natural calamine that is 67.5% ZnO?

chemical compounds - faculty

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