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
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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
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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
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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
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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
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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
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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
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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%
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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
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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
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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
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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
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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
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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
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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
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Dalton’s Atomic Theory – 1803Hypothesis #2
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
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Dalton’s Atomic Theory – 1803Hypothesis #3
3) No atoms are gained/lost/changed inidentity during a chemical reaction,they are just rearranged to producenew substances
Section 12.3
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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
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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
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Electron Shell Distribution
Section 12.4
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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
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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
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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
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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
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Pattern of Ionic Charges
Tend to lose valence electrons Gain electrons invalence shell
Noble Gases
Section 12.4
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Electron Shell DistributionResults in Ionic Charge Pattern
+1 +2 +3 -3 -2 -1
Metals Nonmetals
Section 12.4
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Sodium Ion (Na+) (loses the electron from theouter shell)
Chloride Ion (Cl-) (gains an electron to fill theouter shell)
Section 12.4
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The Formation of Sodium Chloride
NaCl - formula unit of sodium chloride, the smallest combinationof ions that gives the compound formula
Section 12.4
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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
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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+
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Lewis SymbolsFor the First Three Periods of the Representative Elements
Section 12.4
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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
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Electron Shell Distribution
+1 +2 +3 -3 -2 -1
Metals Nonmetals
Section 12.4
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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
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Electron Shell Distribution
+1 +2 +3 -3 -2 -1
Metals Nonmetals
Section 12.4
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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
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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
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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
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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
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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
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Melted Salt (NaCl) Conducts Electricity
Aqueous solutions in which ionic compounds have been dissolvedalso conduct electricity.
Section 12.4
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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
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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
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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
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Covalent Bonding in the H2 molecule
0.074 nm is the distance at which the two H atomsare the most stable
Section 12.5
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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)
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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
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Recall the Lewis Symbols for the FirstThree Periods of Representative Elements
Section 12.5
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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
• Group 5A tend to form 3 bonds
• Group 4A tend to form 4 bonds
Section 12.5
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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
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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
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Double Bond Example – CO2
Section 12.5
..
... .
..• :O. . C . .O:
.. ..• :O : : C : : O :• or
.. ..• : O==O :
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Triple Bond Example – N2
Section 12.5
• :N:::N:
• or
. .
. .• :N. .N:
__________________________________________• :N N:
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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
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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
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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]-
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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
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Comparison of Properties of Ionic andCovalent Compounds
Section 12.5
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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
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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
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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
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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
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Electronegativity Values
Section 12.5
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Summaryof Ionic
andCovalentBonding
Section 12.5
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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
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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
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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
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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
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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
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--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
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Polar and Nonpolar Liquids
A stream ofpolar watermolecules(left) isdeflected by acharge. Astream ofnonpolar CCl4
molecules isnot deflected
Section 12.5
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Types of Molecules with Polar Bonds butNo Resulting Dipole
Section 12.5
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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
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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
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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
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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
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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
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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
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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
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Hydrogen Bonding in Water
The forces ofattraction (reddots) existbetween thehydrogen (+)atom of onemolecule andthe oxygen (-)atom of anothermolecule
Section 12.6
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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
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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
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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
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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