chapter 12 chapter 12: chemical bonding chemical...

Chapter 12

Chemical BondingChapter 12: Chemical Bonding

Homework: All questions on the “Multiple-Choice” and the odd-numbered questions on“Exercises” sections at the end of the chapter.

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Chemical Bonding

• This chapter focuses on chemical bondingand its role in compound formation

• Virtually everything in nature depends onchemical bonds– Proteins, carbohydrates, and fats that make up

living matter are complex molecules held bychemical bonds

• Rock/minerals on earth – compounds heldtogether by chemical bonds

• Chemical bonding results fromelectromagnetic forces between the electronsand nuclei

Intro


Law of Conservation of Mass

• No detectable change in the total massoccurs during a chemical reaction

• If the total mass involved in a chemicalreaction is precisely measured beforeand after the reaction, there is nodifference

• Discovered in 1774 by FrenchmanAntoine Lavoisier

Section 12.1


Law of Conservation of Mass

• If a candle is burned in an airtight container ofoxygen there is no detectable change in themass

Section 12.1


Law of Conservation of Mass - Example

• The complete burning in oxygen (O) of4.09 g of carbon (C) produces 15.00 g ofcarbon dioxide (CO2). How many gramsof oxygen reacted?

• Carbon + oxygen carbon dioxide

• 4.09 g + ? 15.00 g

• Obviously the “?” = 10.91 g

• \ of the 15.00 g of CO2

• 4.09 g = C & 10.91 g = O

Section 12.1


Formula Mass

• Recall that the atomic mass (AM) of anelement is the average mass of all itsnaturally occurring isotopes

– Round off these values to the nearest 0.1 u

• The formula mass (FM) of a compound isthe sum of the atomic masses given in itsformula

• For example: CH4 = 12 u + (4 x 1.0 u) =16 u

Section 12.2


Calculating Formula Masses

• Find the formula mass (FM) of leadchromate, PbCrO4 – used for yellow lineson streets

• Using the Periodic Table, look up theatomic masses of Pb, Cr, and O

• Pb (207.2 u), Cr (52.0 u), O (16.0 u)

• Formula Mass = 207.2 u + 52.0 u + (4 x16 u)

• FM = 323.2 u

Section 12.2


Law of Definite Proportions

• Different samples of a pure compoundalways contain the same elements in thesame proportion by mass.

• For Example:– 9 g H2O = 8 g Oxygen + 1 g Hydrogen

– 18 g H2O = 16 g Oxygen + 2 g Hydrogen

– 36 g H2O = 32 g Oxygen + 4 g Hydrogen

• In each case the ratio (or proportion) bymass of Oxygen to Hydrogen is 8 to 1

Section 12.2


Calculating Percentage by Mass of anElement

Section 12.2

• % X by mass = (atoms of X in formula) x (AMx) X 100

FMcpd• H2O for example

• % 0 by mass = (1) x (16.0 u) X 100 = 88.9%

18.0

• % H by mass = 11.1%


Calculating Percentage by Mass for CO2

• “Dry Ice” is CO2

• AM (atomic mass) of C = 12.0 u & O = 16.0 u

• FM (formula mass) of CO2 =

– 12.0 u + (2 x 16.0 u) = 44.0 u

• % mass of C = (1 x AMc/FMCO2) x 100 = ???%

• % mass of C = (1 x 12.0 u/44.0 u) x 100 = 27.3%

• Since the % mass of C = 27.3%

• \ the % mass of O = 72.7 %

Section 12.2


Calculating Percentage by Mass for Al2O3

• Mineral corundum (ruby & sapphire) is Al2O3

• AM (atomic mass) of Al = 27.0 u & O = 16.0 u

• FM (formula mass) of Al2O3 =

– (2 x 27.0 u) + (3 x 16.0 u) = 102.0 u

• % mass of O = (3 x AMO/FMAl2O3) x 100 = ???%

• % mass of O = (3 x 16.0 u/102.0 u) x 100 =47.1%

• Since the % mass of O = 47.1%

• \ the % mass of Al = 52.9 %

Section 12.2


Definite Proportions

• When a compound is broken down, itselements are found in a definiteproportion by mass

• Also, when the same compound isformed, the elements will combine inthat same proportion by mass

Section 12.2


Limiting & Excess Reactants

• If constituent elements are not mixed inthe correct proportions then

• One of the elements will be usedcompletely up and is called the limitingreactant

• And one of the elements will onlypartially be used up and is called theexcess reactant

• Let’s look at an example …

Section 12.2


Law of Definite ProportionsNote that the Law of Conservation of Mass is also satisfied!

Correct %’s

Excess SLimited Cu

Excess CuLimited S

Section 12.2


Dalton’s Atomic Theory

• In 1803, John Dalton proposed threehypotheses to explain the following twolaws

– Law of Conservation of Mass

– Law of Definite Proportions

Section 12.3


Dalton’s Atomic Theory – 1803Hypothesis #1

1) Each element is composed of smallindivisible particles called atoms

– Atoms are identical for that element, butdifferent from other atoms

Section 12.3



2) Chemical combination is simply thebonding of a definite number of atomsto make one molecule of thecompound

– A given compound always has the samerelative numbers and types of atoms

Section 12.3



3) No atoms are gained/lost/changed inidentity during a chemical reaction,they are just rearranged to producenew substances

Section 12.3


Dalton’s Atomic Theory - 1803

• Over the years more and moresupporting evidence for Dalton’sconcept of the atom has accumulated

• Although there have been manymodifications to his basic ideas, theyhave worked so well and for so long thatwe now call it the atomic theory

• Dalton’s atomic theory is thecornerstone of modern chemistry

Section 12.3


A Little Review

• In Chapter 11 we learned that elements in thesame group have the same # of valenceelectrons.– Similar compounds –> LiCl, NaCl, KCl, RbCl, CsCl

• Because of this behavior, we know that thevalence electrons are the ones involved incompound formation.

• Group 8A are the “noble gases” and generallydo not bond with other atoms.

• Chemists have concluded that having eightelectrons in the outer shell is very stable.

Section 12.4


Electron Shell Distribution

Section 12.4


Octet Rule

1) Valence electrons are the ones involved incompound formation.

2) Eight electrons in the outer shell is verystable.

• The vast majority of compounds can beexplained by combining these twoconclusions.

• In forming compounds, atoms tend to gain,lose, or share valence electrons to achieveelectron configurations with eight electron inthe outer shell. (H is an exception w/ only 2in outer shell.)

Section 12.4


Bonding

• Individual atoms can achieve this “noblegas” electron configuration (8 in outershell) in two ways:

– By transferring (gaining or losing) electrons

– By sharing electrons

• Bonding by transfer of electrons iscalled ionic bonding and will bediscussed in this section.

Section 12.4


Ionic Bonding

• In the transfer of electrons:– One or more atoms lose their valence

electrons

– Another one or more atoms gain thesesame electrons

– In order to achieve noble gas electronconfigurations

• Compounds formed by this electrontransfer process are called ioniccompounds.

Section 12.4


Ions

• An ion is formed due to the loss or gain ofelectrons that destroy the electrical neutralityof the atom and produces a net positive ornegative electric charge.

• The net electric charge on an ion is thenumber of protons minus the number ofelectrons. (p – e = net charge)

• Metals (left side of Periodic Table) tend tolose one or two electrons.

• Nonmetals (right side of Periodic Table) tendto gain electrons.

Section 12.4


Pattern of Ionic Charges

Tend to lose valence electrons Gain electrons invalence shell

Noble Gases

Section 12.4


Electron Shell DistributionResults in Ionic Charge Pattern

+1 +2 +3 -3 -2 -1

Metals Nonmetals

Section 12.4


Sodium Ion (Na+) (loses the electron from theouter shell)

Chloride Ion (Cl-) (gains an electron to fill theouter shell)

Section 12.4


The Formation of Sodium Chloride

NaCl - formula unit of sodium chloride, the smallest combinationof ions that gives the compound formula

Section 12.4


Sodium Chloride (NaCl) – schematicdiagram of a crystal showing a formula unit

• Note that it is actuallyimpossible to associateany one Na+ with onespecific Cl-

• Thus it is somewhatinappropriate to refer to a“molecule” of any ioniccompound.

Section 12.4


Lewis Symbol

• Lewis Symbol – the nucleus and innerelectrons of an atom/ion represented bythe element’s symbol, and the valenceelectrons shown by dots

Section 12.4

: ::• .Cl: or :Cl:—

• Na. or Na+


Lewis SymbolsFor the First Three Periods of the Representative Elements

Section 12.4


Ions

• Cations – positive ions, generally metals

– Elements that tend to lose electrons

– The positive charge will be equal to the number ofvalence electrons in the atom (its group number.)

• Anions – negative ions, generally nonmetals

– Elements that tend to gain electrons

– The negative charge on the nonmetal’s ion will bethe number of valence electrons in the atom (itsgroup number) minus 8.

Section 12.4



+1 +2 +3 -3 -2 -1

Metals Nonmetals

Section 12.4


Pattern of Charges??

• Valence electrons are lost by metals andgained by nonmetals generally to theextent necessary to acquire eightelectrons into the most outer shell, thatis, to acquire an electron configurationisoelectronic with a noble gas. (sameelectron configuration)

• For example: Al3+ is isoelectric with Ne

• or S2- is isoelectronic with Ar.

Section 12.4



+1 +2 +3 -3 -2 -1

Metals Nonmetals

Section 12.4


Ionic Bonds and Compounds

• Ionic bond – electrical forces that hold the ionstogether in the crystal lattice of an ioniccompound

• In every ionic compound, the total charge in theformula adds up to zero and the compoundexhibits electrical neutrality.

• In NaCl the ratio of Na+ to Cl- is always 1 to 1.

• In CaCl2 the ratio of Ca2+ to Cl- is always 1 to 2.

Section 12.4


Formulas for Ionic Compounds

• The numbers of atoms of the variouselements in a compound aredetermined by:

1) The total electrical charge for thecompound is zero

2) All the individual atoms have noble gasconfigurations in their outer shell

Section 12.4


Writing formulas for Ionic Compounds –an Example

• Write the formula for calcium phosphate, themajor component of bones.

• Ca is in Group 2A 2+ ionic charge

• Phosphate is the polyatomic ion (section 11.5)PO4 3- ionic charge

• Therefore when combining these two ionsneutrality can be attained with three Ca2+ andtwo PO4

3-.

• Ca3(PO4)2

Section 12.4


Confidence Exercise

• Write in the formulas for the ionic compounds formedby combining each metal ion (M) with each nonmetalion (X.)

MX

MX2

MX3

M2X

MX

M2X3

M3X

M3X2

MX

Section 12.4


Ionic Compounds

• Due to the very strong forces of attractionbetween oppositely charged ions

– Ionic compounds are always crystalline solids andalso have high melting and boiling points

• Ionic compounds also have a specificproperty when an electric current is passedthrough them.

– Solid ionic compounds DO NOT conduct electricity(because ions cannot move.)

– Melted ionic compounds will CONDUCTelectricity.

Section 12.4


Melted Salt (NaCl) Conducts Electricity

Aqueous solutions in which ionic compounds have been dissolvedalso conduct electricity.

Section 12.4


Stock System –For metals that form two ions

• Stock System – place in parentheses directly afterthe metal’s name a Roman numeral giving the valueof the metal’s ionic charge– CrCl2 chromium(II) chloride (usually blue)– CrCl3 chromium(III) chloride (usually green)

Section 12.4


Stock System: Example

• A certain compound of gold and sulfur has theformula Au2S. What is the Stock systemname?

• Au = either 2+ or 1+ & S = 2-

• Therefore the Stock system name = gold(I)sulfide

• Cu = either 2+ or 1+ & F = 1-

• CuF = copper(I) fluoride

• CuF2 = copper(II) fluoride

• (The old names for these compounds werecuprous fluoride and cupric flouride.)

Section 12.4


Covalent Bonding

• When a pair of electrons is shared by twoatoms, a covalent bond exists between theseatoms

– The two electrons no longer orbit an individualnucleus, but are shared equally by both nuclei

• If the covalent bond is between atoms of thesame element, the molecule formed is that ofan element -- H2 (hydrogen gas)

• Covalent bonds between atoms of differentelements form molecules of compounds – HCl

Section 12.5


Covalent Bonding in the H2 molecule

0.074 nm is the distance at which the two H atomsare the most stable

Section 12.5


Lewis Symbol Use

• Hydrogen gas (H2)

• H. .H H : H or H—H (shows twohydrogen atoms each sharing both valenceelectrons – a covalent bond)

• Hydrogen chloride (hydrochloric acid) (HCl)

Section 12.5

..

.. ..

..

..

..• H. .Cl: H : Cl : or H-Cl : (shows the H

and Cl atoms sharing two electrons – a covalentbond)


Stable Covalent Molecules

• Stable Covalent Molecules form whenthe atoms share electrons in such a wayas to give all atoms a share in a noblegas configuration

• Recall the Octet Rule – in formingcompounds, atoms tend to gain, lose, orshare electrons to achieve electronconfigurations of the noble gases

Section 12.5


Recall the Lewis Symbols for the FirstThree Periods of Representative Elements

Section 12.5


Covalent Bonds & Groups

• Noble Gases (8A) tend to form 0 bonds

• Hydrogen and Group 7A tend to form 1bond

• Group 6A tend to form 2 bonds



Section 12.5


Number of Covalent Bonds expected byCommon Nonmetals

Exceptions are uncommon in Periods 1 & 2, but occur with more frequencystarting with Period 3

Section 12.5


Double/Triple Bonding

• When an element has 2, 3, or 4unpaired valence electrons, its atomswill sometimes share more than one ofthem with another atom

• Double and Triple bonds between twoatoms are possible

Section 12.5


Double Bond Example – CO2

Section 12.5

..

... .

..• :O. . C . .O:

.. ..• :O : : C : : O :• or

.. ..• : O==O :


Triple Bond Example – N2

Section 12.5

• :N:::N:

• or

. .

. .• :N. .N:

__________________________________________• :N N:


Drawing Lewis Structures for SimpleCovalent Compounds – use board

• Chloroform CHCl3• C (four bonds); H (one bond); Cl (one

bond) each

• \ only C can be the central atom

• Hydrogen peroxide H2O2

• H (one bond); O(two bonds)

• \ only O can connect two atoms

Section 12.5


Covalent Bonding –Misc. Info.

• Unlike ionic compounds, covalentcompounds are composed of individualmolecules with a specific molecularformula

• Carbon tetrachloride (CCl4) consists ofmany individual CCl4 molecules

• Within a molecule the covalent bonds arestrong, but the individual molecules onlyweakly attract each other

Section 12.5


Ionic & Covalent??

• Some compounds contain both ionicand covalent bonds – sodium hydroxide(NaOH)

• OH- is a polyatomic ion that has acovalent bond between the O & H

• But there is an ionic bond between theNa+ and the OH-, holding the wholemolecule together

Section 12.5

..

..• Na+[:O:H]-


Rules to Predict Ionic or Covalent??

• Compounds formed of only nonmetals arecovalent (except ammonium compounds)

• Compounds of metals and nonmetals aregenerally ionic

• Compounds of metals with polyatomic ionsare ionic

• Compounds that are gases, liquids, or low-melting-point solids are covalent

• Compounds that conduct electricity whenmelted are ionic

Section 12.5


Comparison of Properties of Ionic andCovalent Compounds

Section 12.5


Predicting Bonding Type – examples

• KF ionic, a metal and nonmetal

• SiH4 covalent, all nonmetals

• Ca(NO3)2 ionic, metal and polyatomicion

• X (a gas at room temp) covalent, agas

• Y (melts at 900oC, then conductselectricity) ionic

Section 12.5


Polar Covalent Bonding

• Remember that in covalent bonding,electrons are shared, but …

– These bonds are not always shared equally

• Unless the atoms are the same element, thebonding electrons spend more time aroundthe more nonmetallic element

– The sharing is unequal

• The is called a polar covalent bond, indicatinga slightly positive end and a slightly negativeend

Section 12.5


Electronegativity

• Electronegativity (EN) – a measure ofthe ability of an atom in a molecule todraw bonding electrons to itself

• Electronegativity also displays definitetrends on the Periodic Table

– Increases across a period

– Decreases down a group

Section 12.5


Electronegativity – an Example

• Consider HCl H (EN=2.1) & Cl (EN=3.0)

– Note that the Cl is more electronegative

• The two bonding electrons tend to spendmore time at the Cl- end \ resulting in a polarcovalent bond

• Polarity can be represented –

• The head of arrow points to the moreelectronegative atom and the other sidemakes a “plus”

Section 12.5


Electronegativity Values

Section 12.5


Summaryof Ionic

andCovalentBonding

Section 12.5


Showing the Polarity of Bonds - example

• Use arrows to show the polarity of thecovalent bonds of H2O

• O (EN=3.5) & H (EN=2.1)

• Cl (EN=3.0) & C (EN=2.5)

Section 12.5


Showing the Polarity of Bonds - example

• Use arrows to show the polarity of thecovalent bonds of CCl4

• O (EN=3.5) & H (EN=2.1)

• Cl (EN=3.0) & C (EN=2.5)

Section 12.5


Polar Bonds and Polar Molecules

• The molecule as a whole, as well asbonds, can have polarity

• A molecule is polar if electrons are moreattracted to one end of the molecule

• Such a molecule has a slightly (-) endand a slightly (+) end

• This type of molecule is said to have adipole or is called a polar molecule

Section 12.5


Polar Bonds & Polar Molecules

H (EN=2.1) Be (EN=1.5)

Nonpolar Molecule

Section 12.5

H (EN=2.1) Cl (EN=3.0)

Polar Molecule


Water Molecule – it is polar!

• BUT• The water molecule is

actually angular (105o)and \ has a positive andnegative end (dipole)

Section 12.5

• If water was a linearmolecule it would benonpolar

• Therefore, in order to determine if a molecule ispolar (dipole), one must know the molecule’sshape


--Summary--Polar Bonds and Polar Molecules

• If the bonds in a molecule are nonpolar,the molecule can only be nonpolar

• A molecule with only one polar bondhas to be polar

• A molecule with more than one polarbond will be nonpolar if the shape of themolecule causes the polarities of thebonds to cancel, otherwise it will bepolar

Section 12.5


Polar and Nonpolar Liquids

A stream ofpolar watermolecules(left) isdeflected by acharge. Astream ofnonpolar CCl4

molecules isnot deflected

Section 12.5


Types of Molecules with Polar Bonds butNo Resulting Dipole

Section 12.5


Dissolving

• Why does water dissolve table salt butdoes not dissolve oil?

• The polar nature of the water moleculecauses them to interact with an ionicsubstance such as salt

– The positive ends of the water moleculesattract negative ions

– The negative ends of the water moleculesattract positive ions

Section 12.6


Polar Water Molecules Dissolve

• If the attraction of the polar watermolecules overcomes the attractionbetween the ions in the crystal, the saltdissolves

• This type of attraction is called an ion-dipole interaction

Section 12.6


Sodium Chloride Dissolving in Water

• The (-) ends of the polar water molecules attract/surround the (+) Na

• The (+) ends of the polar water molecules attract/surround the (-) Cl

Section 12.6


Polar and Nonpolar Substances

• Two polar substances tend to dissolve ineach other

– They are said to have a dipole-dipole interaction

• Two nonpolar substances also tend to mixwell, but not for the same reason

– Nonpolar molecules of two types simply have noaffinity for each other and \ evenly disperse

– Gas and Oil are nonpolar molecules that mix well

Section 12.6


Polar and Nonpolar Substances

• In general like dissolves like

– Polar substances tend to mix well in other polarsubstances

– Nonpolar substance tend to mix well in othernonpolar substances

• Unlike substances do not tend to mix well

– The polar molecules tend to gather together andexclude the nonpolar molecules

– For example oil (nonpolar) does not mix well withwater (polar)

Section 12.6


Hydrogen Bonding

• Hydrogen bond – a special kind ofdipole-dipole interaction

• Hydrogen bonding occurs wheneverhydrogen atoms are covalently bondedto small, highly electronegative atoms

• In general, O, F, and N meet thesecriteria

Section 12.6


Hydrogen Bonding

• When hydrogen is covalently bonded toone of these three atoms …

• The bond is very polar and thehydrogen atom is comparatively small

• Thus the partial positive charge on thehydrogen is highly concentrated

• Resulting in the hydrogen atom havingan electrical attraction for nearby O, F,or N atoms in neighboring molecules

Section 12.6


Hydrogen Bonding in Water

The forces ofattraction (reddots) existbetween thehydrogen (+)atom of onemolecule andthe oxygen (-)atom of anothermolecule

Section 12.6


Hydrogen Bonds

• Hydrogen bonds are strong enough tohave a significant effect on theproperties of the substance

– Hydrogen bonds are about 5-10% thestrength of covalent bonds

Section 12.6


Hydrogen Bonding Affects theProperties of a Substance

• One of the most pronounced effects thatresults from hydrogen bonding is thepredicted change in a substance’s boilingpoint

• Note on the following graph - the threesubstances with hydrogen bonding havesignificantly higher boiling points

– Hydrogen bonding does not occur in CH4, itsboiling point shows a normal pattern relative to theother hydrogen compounds of Group 4A

Section 12.6


Hydrogen Bonding at Work

In general, the boilingpoints of similarcompounds increasewith increasing formulamass. Due tohydrogen bonding theboiling points of H2O,HF, and NH3 are allanomalously high.

Section 12.6


Hydrogen Bonding and Density

• Hydrogen bonding and the shape of the water molecule result inice having a more open structure. In most other substances thesolid phase is more dense than the liquid phase

Water (H2O) Benzene (C6H6)

Section 12.6

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