Chemistry 1AChapter 3

Covalent Bond Formation

Covalent Bond• A link between atoms due to the sharing of

two electrons. This bond forms between atoms of two nonmetallic elements.– If the electrons are shared equally, there is a

even distribution of the negative charge for the electrons in the bond, so there is no partial charges on the atoms. The bond is called a nonpolar covalent bond.

– If one atom in the bond attracts electrons more than the other atom, the electron negative charge shifts to that atom giving it a partial negative charge. The other atom loses negative charge giving it a partial positive charge. The bond is called a polar covalent bond.

Polar Covalent Bond

Ionic Bond• The attraction between cation and anion. • Atoms of nonmetallic elements often attract

electrons so much more strongly than atoms of metallic elements that one or more electrons are transferred from the metallic atom (forming a positively charged particle or cation), to the nonmetallic atom (forming a negatively charged particle or anion).

• For example, an uncharged chlorine atom can pull one electron from an uncharged sodium atom, yielding Cl− and Na+.

Ionic Bond Formation

Sodium Chloride, NaCl, Structure

Electronegativities

Electronegativity, a measure of the electron attracting ability of atoms in chemical bonds is used to predict…

• whether a chemical bond is nonpolar covalent, polar covalent, or ionic.

• which atom in a polar covalent bond is partial negative and which is partial positive.

• which atom in an ionic bond forms the cation and which forms the anion.

• which of two covalent bonds are more polar.

Bond Types

Which atom in a polar covalent bond is partially negative and which is partially positive?

higher electronegativity↓

partial negative charge

lower electronegativity↓

partial positive charge

Which of two bonds is more polar?

The greater the ΔEN is, the more polar the bond.

Types of Compounds

• All nonmetallic atoms usually leads to all covalent bonds, which from molecules. These compounds are called molecular compounds.

• Metal-nonmetal combinations usually lead to ionic bonds and ioniccompounds.

Classification of Compounds

Summary• Nonmetal-nonmetal combinations

(e.g. HCl)– Covalent bonds– Molecules– Molecular Compound

• Metal-nonmetal combinations (e.g. NaCl)– Probably ionic bonds– Alternating cations and anions in

crystal structure– Ionic compound

Valence Electrons

• The valence electrons for each atom are the most important electrons in the formation of chemical bonds.

• The number of valence electrons for the atoms of each element is equal to the element’s A-group number on the periodic table.

• Covalent bonds often form to pair unpaired electrons and give the atoms of the elements other than hydrogen and boron eight valence electrons (an octet of valence electrons).

Valence Electrons and A-Group Numbers

Electron-Dot Symbols and Lewis Structures

• Electron-dot symbols show valence electrons.

• Nonbonding pairs of valence electrons are called lone pairs.

Lewis Structures

• Lewis structures represent molecules using element symbols, lines for bonds, and dots for lone pairs.

Most Common Bonding Patterns for Nonmetals

Element Number of bonds

Number of lone pairs

H 1 0

C 4 0

N, P 3 1

O, S, Se 2 2

F, Cl, Br, I 1 3

Drawing Lewis Structures

• A later chapter describes procedure that allows you to draw Lewis structures for many different molecules.

• Many Lewis structures can be drawn by attempting to give each atom in a molecule its most common bonding pattern.

Lewis Structure for Methane, CH4

• Carbon atoms usually have 4 bonds and no lone pairs.

• Hydrogen atoms have 1 bond and no lone pairs.

Tetrahedral Geometry

Methane, CH4

Lewis Structure for Ammonia, NH3

• Nitrogen atoms usually have 3 bonds and 1 lone pair.

• Hydrogen atoms have 1 bond and no lone pairs.

Ammonia, NH3

Lewis Structure for Water, H2O

• Oxygen atoms usually have 2 bonds and 2 lone pairs.

• Hydrogen atoms have 1 bond and no lone pairs.

Water, H2O

Water Attractions

Liquid Water

Study Sheets for Converting Between Names and Formulas

• For each type of compound,– How recognize from formula– How recognize from name– How write formula– How write name

Binary Covalent

Common Names

–H2O, water–NH3, ammonia–CH4, methane –C2H6, ethane–C3H8, propane

Naming Binary Covalent Compounds

• If the subscript for the first element is greater than one, indicate the subscript with a prefix.– We do not write mono- on the first name.– 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 that begins with a vowel (oxygen or iodine).

• Follow the prefix with the name of the first element in the formula.

Naming Binary Covalent Compounds

• Write a prefix to indicate the subscript for the second element.

• Write the root of the name of the second symbol in the formula.

• Add -ide to the end of the name.

Prefixes

mon(o)ditritetr(a)pent(a)

hex(a)hept(a)oct(a)non(a)dec(a)

Roots of Nonmetals

H hydr-C carb-N nitr-P phosph-O ox-S sulf-Se selen-

F fluor-Cl chlor-Br brom-I iod-

Binary Covalent Names without Prefixes

Formula Complete name Abbreviated Name

HF Hydrogen monofluoride

Hydrogen fluoride

HCl Hydrogen monochloride

Hydrogen chloride

HBr Hydrogen monobromide

Hydrogen bromide

HI Hydrogen moniodide Hydrogen iodide

H2S Dihydrogen monosulfide

Hydrogen sulfide

The Making of an Anion

The Making of a Cation

Monatomic Ions

Sodium Chloride, NaCl, Structure

Monatomic Ion Names

• Monatomic Cations– (name of metal)

• Groups 1, 2, and 3 metals• Al3+, Zn2+, Cd2+, Ag+

– (name of metal)(Roman numeral)• All metallic cations not mentioned above

• Monatomic Anions– (root of nonmetal name)ide

Naming Cations with Two Possible Charges

Ion Systematic Name

Nonsystematic Name

Fe2+ iron(II) ferrous

Fe3+ iron(III) ferric

Cu+ copper(I) cuprous

Cu2+ copper(II) cupric

Hydride H−

Nitride N3−

Phosphide P3−

Oxide O2−

Sulfide S2−

selenide Se2−

fluoride F−

chloride Cl−

bromide Br−

iodide I−

Monatomic Anions

CsCl and NH4Cl structure

Polyatomic Ions

Ion Name Ion Name

NH4+ ammonium NO3

− nitrate

OH− hydroxide SO42− sulfate

CO32− carbonate C2H3O2

− acetate

PO43− phosphate

(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

Convention for Naming Oxyanions

Relationship General Name

Example Name

Example Formula

One more O per(root)ate perchlorate ClO4−

(root)ate chlorate ClO3−

One less O (root)ite chlorite ClO2−

Two less O hypo(root)ite hypochlorite ClO−

Polyatomic Ions with Hydrogen

• HCO3− hydrogen carbonate

• HSO4− hydrogen sulfate

• HSO3− hydrogen sulfite

• HS− hydrogen sulfide• HPO4

2− hydrogen phosphate• H2PO4

− dihydrogen phosphate

Systematic and Nonsystematic Names

Formula Systematic Name Nonsystematic Name

HCO3− hydrogen carbonate bicarbonate

HSO4− hydrogen sulfate bisulfate

HSO3− hydrogen sulfite bisulfite

Study Sheets for Converting Between Names and Formulas

• For each type of compound,– How recognize from formula– How recognize from name– How write formula– How write name

Names and Formulas of Ionic Compounds

• Name– (name of cation) (name of anion)

• Formula Steps– Determine the Formula, including

charge, for the cation and anion. – Determine the ratio of the ions that

yields zero overall charge.

Cation Names

Metals with one possible charge (Al,

Zn, Cd, and Groups 1, 2, 3)

name of metal

Metals with more than one possible charge

(the rest)

name(Roman numeral)

polyatomic cations (e.g. ammonium)

name of polyatomic ion

Anion Names

monatomic anion (root of nonmetal name)ide

polyatomic anion name of polyatomic ion

Binary Acid Names and Formulas

• General – HX(aq) hydro(root)ic acid– HF(aq) hydrofluoric acid– HCl(aq) hydrochloric acid– HBr(aq) hydrobromic acid– HI(aq) hydroiodic acid– H2S(aq) hydrosulfuric acid

Oxyacid Names and Formulas

• When enough H+ ions are added to a (root)ate polyatomic ion to neutralize its charge, the (root)ic acid.– nitrate, NO3

−, forms nitric acid, HNO3

– sulfate, SO42−, forms sulfuric acid, H2SO4

– phosphate, PO43−, forms phosphoric acid, H3PO4

– carbonate, CO32−, forms carbonic acid, H2CO3

– chlorate, ClO3−, forms chloric acid, HClO3

– bromate, BrO3−, forms bromic acid, HBrO3

– iodate, IO3−, forms iodic acid, HIO3

Oxyacid Names and Formulas (cont)

Relationship General Name

Example Name

Example Formula

One more O per(root)ic acid perchloric acid HClO4

(root)ic acid chloric acid HClO3

One less O (root)ous acid chlorous acid HClO2

Two less O hypo(root)ous acid

hypochlorous acid HClO

Common Alcohols and Sugars

• CH3OH methanol (methyl alcohol)

• C2H5OH ethanol (ethyl alcohol)

• C3H7OH 2-propanol (isopropyl alcohol)

• C6H12O6 glucose (or fructose or galactose)

• C12H22O11 sucrose(or maltose or lactose)

Condensation (Gas to Liquid)

Types of Particles

Type of substance

Particles to visualize

Type of substance

Particles to visualize

metalscations in a

sea of electrons

ionic cations and anions

noble gases atoms molecular molecules

carbon (diamond) atoms

other nonmetallic elements

molecules

Types of Attractions –Metallic Elements

• Metallic bonds = the attraction between the positive metal cations and the negative electrons that surround them.

Types of Attractions –Ionic Compounds

• Cations and Anions held together by ionic bonds

Types of Attractions –Molecular Compounds

• Nonpolar molecular compounds (e.g. hydrocarbons, CaHb) – London forces

• Polar molecular compounds without H-F, O-H, or N-H (e.g. HCl) – dipole-dipole attractions

• Polar molecular compounds with H-F, O-H, or N-H (e.g. HF, H2O, NH3, alcohols, and sugars) – hydrogen bonds

Dipole-Dipole Attractions

Dipole-Dipole Attractions in a Liquid

Hydrogen Bonds

• Hydrogen bond = the intermolecular attraction between the partially negative O, N, or F of one molecule and a partially positive H connected to O, N, or F of another molecule.

• Relatively strong

Hydrogen Bonds in HFIn HF, the hydrogen bond is between the partial positive H of one HF molecule and the partial negative F of another HF molecule.

Hydrogen Bonds in WaterIn H2O, the hydrogen bond is between the partial positive H of one H2O molecule and the partial negative O of another H2O molecule.

Hydrogen Bonds in MethanolIn CH3OH, the hydrogen bond is between the partial positive H of one CH3OH molecule and the partial negative O of another CH3OH molecule.

Hydrogen Bonds in AmmoniaIn NH3, the hydrogen bond is between the partial positive H of one NH3 molecule and the partial negative N of another NH3molecule.

London Forces

London Forces in Polar Molecules

Why Larger Molecules Have Stronger London Forces

Types of Attractions –Carbon

• Diamond - Carbons atoms held together by covalent bonds, forming huge 3-dimensional molecules.

• Graphite - Carbons atoms held together by covalent bonds, forming huge 2-dimensional molecules held together by London forces.

• Fullerenes - Carbons atoms held together by covalent bonds, forming 3-dimensional molecules held together by London forces.

• See the following website.

http://preparatorychemistry.com/Bishop_Jmol_Carbon.htm

Predicting Types of Attractions

Types of Particles and Attractions - ElementsType of element

Particles to visualize

Examples Type of attraction

metals cations in a sea of electrons gold, Au metallic

bonds

noble gases atoms xenon, Xe London forces

carbon (diamond) atoms C(dia) covalent

bonds

other nonmetallic elements

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

London forces

Types of Particles and Attractions - Compounds

Type of compound Particles to visualize

Examples Type of attraction

ionic cations and anions NaCl ionic

bonds

nonpolar molecular molecules hydrocarbons London forces

polar molecular w/out H-F, O-H, or N-H molecules HCl dipole-

dipole

polar molecular with H-F, O-H, or N-H molecules HF, H2O, NH3

alcohols, hydrogen

bonds

Making Phosphoric Acid

• Furnace Process for making H3PO4 to be used to make fertilizers, detergents, and pharmaceuticals. – React phosphate rock with sand and coke at 2000

°C.2Ca3(PO4)2 + 6SiO2 + 10C

→ 4P + 10CO + 6CaSiO3– React phosphorus with oxygen to get

tetraphosphorus decoxide.4P + 5O2 → P4O10

– React tetraphosphorus decoxide with water to make phosphoric acid.

P4O10 + 6H2O → 4H3PO4

Sample Calculations (1)

• What is the maximum mass of P4O10 that can be formed from 1.09 × 104 kg P?

• The formula for P4O10 provides us with a conversion factor that converts from units of P to units of P4O10.

Sample Calculations (2)• What is the minimum mass of water that

must be added to 2.50 × 104 kg P4O10 to form phosphoric acid in the following reaction?

P4O10 + 6H2O → 4H3PO4• The coefficients in the balanced

equation provide us with a conversion factor that converts from units of P4O10to units of H2O.

Goal: To develop conversion factors that will convert between a measurable property (mass) and number of particles

Measurable Property 1↓

Number of Particles 1↓

Number of Particles 2↓

Measurable Property 2

Mass 1↓

Number of Particles 1↓

Number of Particles 2↓

Mass 2

Counting by Weighing for Nails

• Step 1: Choose an easily measurable property.– Mass for nails

• Step 2: Choose a convenient unit for measurement.– Pounds for nails

Counting by Weighing for Nails (cont)

• Step 3: If the measurable property is mass, determine the mass of the individual objects being measured.– Weigh 100 nails: 82 are 3.80 g, 14 are 3.70

g, and 4 are 3.60 g

• Step 4: If the objects do not all have the same mass, determine the weighted average mass of the objects.

0.82(3.80 g) + 0.14(3.70 g) + 0.04(3.60 g) = 3.78 g

Counting by Weighing for Nails (cont)

• Step 5: Use the conversion factor from the weighted average to make conversions between mass and number of objects.

Counting by Weighing for Nails (cont)

• Step 6: Describe the number of objects in terms of a collective unit such as a dozen, a gross, or a ream.

Counting by Weighing for Carbon Atoms

• Step 1: Choose an easily measurable property.– Mass for carbon atoms

• Step 2: Choose a convenient unit for measurement.– Atomic mass units (u) for carbon atoms– Atomic mass unit (u) = 1/12 the mass of a

carbon-12 atom (with 6 p, 6 n, and 6 e−)

Counting by Weighing for Carbon Atoms (cont)

• Step 3: If the measurable property is mass, determine the mass of the individual objects being measured.– For carbon: 98.90% are 12 u and 1.10% are

13.003355

• Step 4: If the objects do not all have the same mass, determine the weighted average mass of the objects.

0.9890(12 u) + 0.0110(13.003355 u) = 12.011 u

Mass Spectrometer

Mass Spectrum

Counting by Weighing for Carbon Atoms (cont)

• Step 5: Describe the number of objects in terms of a collective unit such as a dozen, a gross, or a ream, and use this and the weighted average to create a conversion factor to make conversions between mass and number of objects.

Mole• A mole (mol) is an amount of substance

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

• To four significant figures, there are 6.022 × 1023 atoms in 12 g of carbon-12.

• Thus a mole of natural carbon is the amount of carbon that contains 6.022 × 1023 carbon atoms.

• The number 6.022 × 1023 is often called Avogadro’s number.

Avogadro’s Number

Molar Mass Development

Molar Mass For Elements

• Atomic Mass from the Periodic Table

⎛ ⎞⎜ ⎟⎝ ⎠

(atomic mass) g element1 mol element

Molar Mass Calculation for Carbon

Goal: To develop conversion factors that will convert between a measurable property (mass) and number of particles

Measurable Property 1↓

Number of Particles 1↓

Number of Particles 2↓

Measurable Property 2

Mass 1↓

Moles 1↓

Moles 2↓

Mass 2

Molecular Mass

• Whole = sum of parts• mass of a molecule = sum of the

masses of the atoms in the molecule • molecular mass = the sum of the

atomic masses of the atoms in the molecule

Molar Mass For Molecular Compounds

• Molecular Mass = Sum of the atomic masses of atoms in one molecule

⎛ ⎞⎜ ⎟⎝ ⎠

(molecular mass) g molecular compound1 mol molecular compound

Formula Units

• A formula unit of a substance is the group represented by the substance’s chemical formula, that is, a group containing the kinds and numbers of atoms or ions listed in the chemical formula.

• Formula unit is a general term that can be used in reference to elements, molecular compounds, or ionic compounds.

Formula Unit Examples

Formula Mass for Ionic Compounds

• Whole = sum of parts• Mass of a formula unit = sum of the

masses of the atoms in the formula unit • Formula mass = the sum of the atomic

masses of the atoms in the formula

Molar Mass For Ionic Compounds

• Formula Mass = Sum of the atomic masses of the atoms in a formula unit

⎛ ⎞⎜ ⎟⎝ ⎠

(formula mass) g ionic compound1 mol ionic compound

Molar Mass Development

General Conversions

Units of One Substance to Units of Another

Study Sheets

• Write a description of the “tip-off” that helps you to recognize the type of problem the calculation represents.

• Write a description of the general procedure involved in the particular type of problem.

• Write an example of the type of calculation.

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

• 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.

Sample Study Sheet (2)

• General Steps– Convert the given unit to moles of the

first substance.– Convert moles of the first substance

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

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

Units of Element

to Units of Compound

Beginning and End