unit 8 - bonding organic - dr. g's...

Bonding & Formulas

Unit 8 Organic Chemistry

Unit OverviewTopic 1 - Bonding

Topic 2 - Organic Chemistry

Topic 3 - Polar & Non-Polar Bonds

Topic 4 - Polar & Non-Polar Molecules

Topic 5 - Intermolecular Forces

Topic 6 - Molecular Math

Bonding & Formulas

• What do you already know about BONDING?

!

• If ionic bonds are metal + nonmetal and covalent bonds are nonmetal + nonmetal, what must the third type of bonding be?

• If metals want to lose electrons, how do you think they will bond together?

• If metals conduct electricity, what 2 conditions MUST exist inside all metals?

metal + metal

metal cations with a ‘sea of electrons’

charged particles, mobile

Topic 1

Bonding Comparison

Ionic Covalent Metallicmetal & nonmetal

e- transferred from metal to nonmetal

VERY strong bond

hard, brittle, high boiling / melting points

conduct electricity as liquids or aqueous (NOT at solids

When writing formulas, make sure charges add to

zero. Naming uses ‘ide’,

sometimes roman numerals, Table E.

nonmetal & nonmetal e- shared between

nonmetals weaker bond

soft, usually gas or liquid at room temperature, low boiling / melting pts.

do not conduct electricity in any phase

Names uses ‘ide’ ending with prefixes (mono, di, tri,

tetra)

metal & metal metals loosely hold

electrons so there is a ‘sea of electrons’

can be hard or soft, with varying melting / boiling

points conduct electricity

Naming uses the elemental metal (i.e. Gold, Copper,

Tin, Iron)

Next think about the bonds that form between carbon and chlorineand between aluminum and chlorine. The electronegativity differencebetween C and Cl is 0.6. These two elements form a polar covalent bond.The electronegativity difference between Al and Cl is 1.6. These two ele-ments also form a polar covalent bond. However, the larger differencebetween Al and Cl means that the bond between these two elements ismore polar, with greater partial charges, than the bond between C and Cl is.

Properties of Substances Depend on Bond TypeThe type of bond that forms determines the physical and chemical propertiesof the substance. For example, metals, such as potassium, are very goodelectric conductors in the solid state. This property is the result of metallicbonding. Metallic bonds are the result of the attraction between the electronsin the outermost energy level of each metal atom and all of the otheratoms in the solid metal. The metal atoms are held in the solid because allof the valence electrons are attracted to all of the atoms in the solid.Thesevalence electrons can move easily from one atom to another.They are freeto roam around in the solid and can conduct an electric current.

Table 3 Properties of Substances with Metallic, Ionic, and Covalent Bonds

Covalent Compounds 197

Bond type Metallic Ionic Covalent

Example substance potassium potassium chloride chlorine

Melting point (°C) 63 770 –101

Boiling point (°C) 760 1500 (sublimes) –34.6

Properties • soft, silvery, solid • crystalline, white solid • greenish yellow gas• conductor as a solid • conductor when • not a good conductor

dissolved in water

+ −

Copyright © by Holt, Rinehart and Winston. All rights reserved.

What is Organic Chemistry?

• Atoms of Carbon, Hydrogen, Oxygen and Nitrogen bonding together to form molecules

• Atomic structure and sharing of electrons to form covalent bonds.

Topic 2

Organic hydrocarbons• Watch the following video on

petroleum refining and answer the questions in your booklet.

• http://www.youtube.com/watch?v=jk0WrtA8_T8

http://www.youtube.com/watch?v=jk0WrtA8_T8

More on Petroleum

... where does our oil come

Organic Hydrocarbons

• Hydrocarbons

• “Organic” molecules (making up living things)

• Made of carbon and hydrogen ONLY

• Carbon has ____ valence electron(s), so it will always make ____ bond(s).

• Hydrogen has ____ valence electron(s), so it will always make ____ bond(s).

44

11

Saturated hydrocarbons

• Alkanes

• Hydrocarbons with all __________ bonds.

• Can’t add any more H atoms to it, it’s full. We call that __________________.

• Name alkanes using Table P (prefixes) and Table Q (general formula).

• General formula for alkanes is CnH2n+2

• Examples...

single

saturated

Isomers (of Alkanes)

• Isomer: molecules with the same formula, but different structures.

• Isomers are different molecules with different chemical properties.

• When asked to draw isomers, make sure you include the ‘original’ molecule

• The more carbon atoms in a molecule, the more isomers it has.

Naming Rules

• Name based on longest carbon chain

• Branches get ‘-yl’ ending

• ex. 1 carbon branch = ‘methyl’, 2 carbon branch = ‘ethyl’

• Numbers tell you which carbon(s) the branch(es) are attached to.

• Count carbons so that the numbers in the name are the smallest #s possible.

• A bend in the chain is NOT a branch.

Unsaturated Hydrocarbons

• Unsaturated = can add more hydrogen by breaking double or triple bonds. (Ever heard of unsaturated fats and oils???)

• Alkenes

• Have a double bond. (Represented by a double line)

• General formula is CnH2n

• Example: C2H4

Unsaturated Hydrocarbons• Alkynes

• Have a triple bond. (Represented by a triple line) • General formula is CnH2n-2 • Example: C2H2

!

• Naming Rules: • numbers tell you which carbon atom has the

double / triple bond. • count carbons from end that keeps numbers

smallest.

Polar and Non-Polar Bonds

• Demonstration - ‘Who has more pull?’

• Covalent Bonds: SHARING of electrons

• Up to now we’ve been talking about bonds as being ionic or covalent

• Is Sharing always equal??

• Polar vs. Non-Polar

Topic 3

Covalent BondsNon-Polar Covalent Bond Polar Covalent Bond

electrons are evenly shared between the two atoms

electronegativities of the two atoms are very similar

ALL diatomic substances (made up of two of the same atoms) are non-polar because atoms of the same element will have identical electronegativies.

One atom has more electronegativity

Polar covalent bond has a negative and positive side.

Sigma (+ or -) symbol or by an arrow starting with a + sign. The arrow points to the stronger ‘puller’.

water has a polar covalent bond between the H and O

The Bonding ‘continuum’

• The three types of bonds should be viewed as a continuum..

• no difference in electronegativity = the bond is non-polar

• As the difference increases, the bond becomes more polar...

• until the extreme case, forming an ionic bond!

write this in your reference table!!!

4.00 0.002.0 0.4

Non-Polar Covalent Bond

Polar Molecules Have Positive and Negative EndsHydrogen fluoride, HF, in solution is used to etch glass, such as the vaseshown in Figure 8. The difference between the electronegativity values ofhydrogen and fluorine shows that H and F atoms form a polar covalentbond. The word polar suggests that this bond has ends that are in someway opposite one another, like the two poles of a planet, a magnet, or abattery. In fact, the ends of the HF molecule have opposite partialcharges.

The electronegativity of fluorine (4.0) is much higher than that ofhydrogen (2.2). Therefore, the shared electrons are more likely to befound nearer to the fluorine atom. For this reason, the fluorine atom inthe HF molecule has a partial negative charge. In contrast, the sharedelectrons are less likely to be found nearer to the hydrogen atom. As aresult, the hydrogen atom in the HF molecule has a partial positivecharge. A molecule in which one end has a partial positive charge and theother end has a partial negative charge is called a The HF moleculeis a dipole.

To emphasize the dipole nature of the HF molecule, the formula canbe written as Hδ+Fδ−. The symbol δ is a lowercase Greek delta, which isused in science and math to mean partial. With polar molecules, such asHF, the symbol δ+ is used to show a partial positive charge on one end ofthe molecule. Likewise, the symbol δ− is used to show a partial negativecharge on the other end.

Although δ+ means a positive charge, and δ− means a negative charge,these symbols do not mean that the bond between hydrogen and fluorineis ionic. An electron is not transferred completely from hydrogen to fluo-rine, as in an ionic bond. Instead, the atoms share a pair of electrons,which makes the bond covalent. However, the shared pair of electrons ismore likely to be found nearer to the fluorine atom. This unequal distri-bution of charge makes the bond polar covalent.

dipole.


Figure 7Electronegativity differencescan be used to predict theproperties of a bond. Notethat there are no distinctboundaries between thebond types—the distinctionis arbitrary.

Even electron distribution Uneven electron distribution Separate electron clouds

Electronegativity Difference

Nonpolarcovalent Polar covalent Ionic

0 0.5 2.1

+ –!+ !–

3.3

dipole

a molecule or part of a moleculethat contains both positively andnegatively charged regions

Figure 8Hydrogen fluoride, HF, is an acid that is used to etchbeautiful patterns in glass.


Polar Covalent Bond





dipole.






0 0.5 2.1

+ –!+ !–

3.3

dipole




Ionic Bond





dipole.






0 0.5 2.1

+ –!+ !–

3.3

dipole




Difference in Electronegativity

Practice• Identify the bonds between each pair of elements

as non-polar covalent, polar covalent, or ionic

H and Br C and O

K and Cl Br and Br

Polar Covalent Polar Covalent

Ionic Non-Polar Covalent

More Practice

• Place the following covalent bonds in order from least to most polar.

• A. H–Cl

• B. H–Br

• C. H–S

• D. H–C

Polar & non-Polar Molecules

• Which of these pictures show symmetry?

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Now divide all the symmetrical pictures into 2 groups...

Now we’ll look at the symmetry of molecules.

Topic 4

• They are all _________________ because they consist of only nonmetals.

• That’s why they are called MOLECULES


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Notice how some of them are symmetrical all the way around the molecule and some are not.

What do you notice about all of the shapes that are asymmetrical? What do they all have in common?

Covalent Molecules• To determine if a MOLECULE is polar or nonpolar,

you MUST look at its SHAPE!!

Nonpolar Molecules Polar Molecules

molecules with a SYMMETRICAL SHAPE molecules with an ASYMMETRICAL SHAPE

even distribution of electrons all around the molecule

uneven distribution of electrons all around the molecule

electrons ‘pulled’ to one end of the molecule, which creates a negative ‘pole’ on the

molecule. (The other end then becomes the ‘positive’ pole.

polar asymmetric molecules are also called DIPOLES (meaning ‘two poles’)

Revelation!!

• Now label all of the molecules below as ‘polar’ or ‘nonpolar’.


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Practice: Polar or Not??

• For the molecules below, determine 1) if they have polar or nonpolar bonds in them and 2) if they are a polar molecule or a nonpolar molecule.

NH3 CO2 CCl4 H2O PH3

Regents Practice

Name: ____________________________________________1) Which formula represents a nonpolar molecule?

1) HCl 2) NH3 3) H2S 4) CH42) Which of the following statements explains why low temperature and high pressure are required

to liquefy chlorine gas?1) Chlorine molecules have weak covalent bonds.2) Chlorine molecules have strong intermolecular forces of attraction.3) Chlorine molecules have weak intermolecular forces of attraction.4) Chlorine molecules have strong covalent bonds.

3) What occurs when an atom of chlorine and an atom of hydrogen become a molecule of hydrogenchloride?1) A chemical bond is broken and energy is released.2) A chemical bond is formed and energy is released.3) A chemical bond is formed and energy is absorbed.4) A chemical bond is broken and energy is absorbed.

4) Based on bond type, which of the following compounds has the highest melting point?1) CCl4 2) C6H14 3) CH3OH 4) CaCl2

5) Which type of bond results when one or more valence electrons are transferred from one atom toanother?1) a polar covalent bond2) a nonpolar covalent bond

3) an ionic bond4) a hydrogen bond

6) Which of the following solids has the highest melting point?1) CO2(s) 2) SO2(s) 3) Na2O(s) 4) H2O(s)

7) Which compound contains ionic bonds?1) NO 2) CO2 3) CaO 4) NO2

8) Which bond is least polar?1) As‡‡Cl 2) N‡‡Cl 3) P‡‡Cl 4) Bi‡‡Cl

9) Which type of chemical bond is formed between two atoms of bromine?1) hydrogen 2) ionic 3) metallic 4) covalent

10) Given the Lewis electron-dot diagram:

Which electrons are represented by all of the dots?1) the carbon valence electrons, only2) all of the carbon and hydrogen electrons3) the carbon and hydrogen valence electrons4) the hydrogen valence electrons, only

Questions 11 and 12 refer to the following:

Element X is a solid metal that reacts with chlorine to form a water-soluble binary compound.

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Questions 41 through 44 refer to the following:

Each molecule listed below is formed by sharing electrons between atoms when the atoms within themolecule are bonded together.

Molecule A: Cl2Molecule B: CCl4Molecule C: NH3

41) Explain how the bonding in KCl is different from the bonding in molecules A, B, and C.

42) Draw the electron-dot (Lewis) structure for the NH3 molecule in the box below.

43) Explain why CCl4 is classified as a nonpolar molecule.

44) Explain why NH3 has stronger intermolecular forces of attraction than Cl2.

45) Which type of molecule is CF4?1) nonpolar, with an asymmetrical distribution of charge2) nonpolar, with a symmetrical distribution of charge3) polar, with a symmetrical distribution of charge4) polar, with an asymmetrical distribution of charge

46) Which intermolecular force of attraction accounts for the relatively high boiling point of water?1) metallic bonding2) ionic bonding

3) hydrogen bonding4) covalent bonding

47) Which compound contains both ionic and covalent bonds?1) MgF2 2) CH2O 3) CaCO3 4) PCl3

48) Which Lewis electron-dot diagram is correct for CO2?

1) 2) 3) 4)

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Intermolecular Forces (IMFs)

• the STRONGER the attractions between the particles in a substance, the HIGHER the melting point temperature (or boiling point) becomes.

• Using Table H, which of the 4 substances consists of particles with the strongest forces of attraction between them? ____________________

• How did you decide?

• Do you remember how this relates to vapor pressure?

Topic 5

Ionic Bond

• Ionic substances have higher forces of attraction than covalent substances. The smaller the ions, the larger the attraction.

• If the ions are large, the distances between them are larger and the forces are weaker.

• Attraction also depends on the amount of charge. The more charge the more attraction (CaF2 > NaCl)

• Let’s look at the following table ...

Ionic Bonds

States of Matter and Intermolecular Forces 385

OBJECTIVES

Contrast ionic and molecular substances in terms of their physicalcharacteristics and the types of forces that govern their behavior.

Describe dipole-dipole forces.

Explain how a hydrogen bond is different from other dipole-dipoleforces and how it is responsible for many of water’s properties.

Describe London dispersion forces, and relate their strength to othertypes of attractions.

Comparing Ionic and Covalent CompoundsParticles attract each other, so it takes energy to overcome the forcesholding them together. If it takes high energy to separate the particles ofa substance, then it takes high energy to cause that substance to go fromthe liquid to the gaseous state. The boiling point of a substance is a goodmeasure of the strength of the forces that hold the particles together.Melting point also relates to attractive forces between particles. Mostcovalent compounds melt at lower temperatures than ionic compoundsdo. As shown in Table 1, ionic substances with small ions tend to be solidsthat have high melting points, and covalent substances tend to be gasesand liquids or solids that have low melting points.

KEY TERMS

• intermolecular forces

• dipole-dipole forces

• hydrogen bond

• London dispersion force

11

2

3

4

Intermolecular Forces2S E C T I O N

Table 1 Comparing Ionic and Molecular Substances

State at room Melting point Boiling pointType of substance Common use temperature (°C) (°C)

Ionic substances

Potassium chloride, KCl salt substitute solid 770 sublimes at 1500

Sodium chloride, NaCl table salt solid 801 1413

Calcium fluoride, CaF2 water fluoridation solid 1423 2500

Covalent substances

Methane, CH4 natural gas gas −182 −164

Ethyl acetate, CH3COOCH2CH3 fingernail polish liquid −84 77

Water, H2O (many) liquid 0 100

Heptadecane, C17H36 wax candles solid 22 302


van der Waal Forces• Weak, Short-range, and decrease rapidly as

molecules get farther apart. (don't affect gases)

• These apply to non-polar COVALENT compounds.

• As mass increases, this force increases

390 Chapter 11

Table 4 Boiling Points of the Noble Gases

Substance He Ne Ar Kr Xe Rn

Boiling point (°C) −269 −246 −186 −152 −107 −62

Number of electrons 2 10 18 36 54 86

Figure 12The nonpolar molecules ingasoline are held togetherby London dispersion forces,so it is not a gas at roomtemperature.

London dispersion force

the intermolecular attractionresulting from the uneven distribution of electrons and thecreation of temporary dipoles

+

+

+

—

——

++

+

+

— —

—

—

—

—

+

+

Figure 13Temporary dipoles in molecules cause forces of attraction between the molecules.

a Nonpolar molecules canbecome momentarily polar.

b The instantaneous dipolesthat form cause adjacentmolecules to polarize.

c These London dispersionforces cause the moleculesto attract each other.

London Dispersion ForcesSome compounds are ionic, and forces of attraction between ions ofopposite charge cause the ions to stick together. Some molecules arepolar, and dipole-dipole forces hold polar molecules together. But whatforces of attraction hold together nonpolar molecules and atoms? Forexample, gasoline, shown in Figure 12, contains nonpolar octane, C8H18,and is a liquid at room temperature.Why isn’t octane a gas? Clearly, somesort of intermolecular force allows gasoline to be a liquid.

In 1930, the German chemist Fritz W. London came up with an expla-nation. Nonpolar molecules experience a special form of dipole-dipoleforce called In dipole-dipole forces, the negativepart of one molecule attracts the positive region of a neighboring mol-ecule. However, in London dispersion forces, there is no special part ofthe molecule that is always positive or negative.

London Dispersion Forces Exist Between Nonpolar MoleculesIn general, the strength of London dispersion forces between nonpolarparticles increases as the molar mass of the particles increases. This isbecause generally, as molar mass increases, so does the number of elec-trons in a molecule. Consider the boiling point of the noble gases, asshown in Table 4. Generally, as boiling point increases, so does the num-ber of electrons in the atoms. For groups of similar atoms and molecules,such as the noble gases or hydrogen halides, London dispersion forces areroughly proportional to the number of electrons present.

London dispersion force.


Dipole-Dipole

• The positive end of one molecule attracts the negative end of a neighboring molecule.

• Remember, bonds are polar because of electronegativity (...man).

• The boiling point of a substance tells you about the forces between molecules. The more polar the molecules are = the stronger the dipole-dipole forces = the higher the boiling point.

• Polar asymmetric, middle strength, permanent.

Boiling point State at roomSubstance (°C) Polarity temperature Structure

1-propanol, 97.4 polar liquidC3H7OH

1-propanethiol, 67.8 less liquidC3H7SH polar

Butane, C4H10 −0.5 nonpolar gas

Water, H2O 100.0 polar liquid

Hydrogen −60.7 less gassulfide, H2S polar

Ammonia, −33.35 polar gasNH3

Phosphine, −87.7 less gasPH3 polar

H

CH C

H

H H

C O H

H

H

H

CH C

H

H H

C S H

H

H

Hydrogen BondsCompare the boiling points of H2O and H2S, shown in Table 2. These mol-ecules have similar sizes and shapes. However, the boiling point of H2O ismuch higher than that of H2S. A similar comparison of NH3 with PH3 canbe made.The greater the polarity of a molecule, the higher the boiling pointis. However, when hydrogen atoms are bonded to very electronegativeatoms, the effect is even more noticeable.

Compare the boiling points and electronegativity differences of thehydrogen halides, shown in Table 3. As the electronegativity differenceincreases, the boiling point increases.The boiling points increase somewhatfrom HCl to HBr to HI but increase a lot more for HF. What accounts forthis jump? The answer has to do with a special form of dipole-dipoleforces, called a hydrogen bond.


H

CH HC

H

H H

H

CC

H

H H

Table 3 Boiling Points of the Hydrogen Halides

Substance HF HCl HBr HI

Boiling point (°C) 20 −85 −67 −35

Electronegativity difference 1.8 1.0 0.8 0.5

Table 2 Comparing Dipole-Dipole Forces

hydrogen bond

the intermolecular force occurring when a hydrogen atom that is bonded to a highlyelectronegative atom of one molecule is attracted to twounshared electrons of anothermolecule

www.scilinks.orgTopic: Hydrogen BondingSciLinks code: HW4069

HHO

HHS

H

HH N

H

HH P


Hydrogen ‘Bond’• A special kind of dipole-dipole force.

• Form with a hydrogen atom that is covalently bonded to very electronegative atoms. Hold the F-O-N!!

• The partial positive charge (on Hydrogen) is attracted to the unshared pairs of electrons of neighboring molecules.

• These are strong dipole-dipole forces due to differences in electronegativity AND because hydrogen is so small.

Hydrogen ‘bond’ examples

• The energy of hydrogen bonds is lower than that of normal chemical bonds, but can be stronger than that of other intermolecular forces.

• Can account for many properties of substances.

Boiling point State at roomSubstance (°C) Polarity temperature Structure

1-propanol, 97.4 polar liquidC3H7OH

1-propanethiol, 67.8 less liquidC3H7SH polar

Butane, C4H10 −0.5 nonpolar gas

Water, H2O 100.0 polar liquid

Hydrogen −60.7 less gassulfide, H2S polar

Ammonia, −33.35 polar gasNH3

Phosphine, −87.7 less gasPH3 polar

H

CH C

H

H H

C O H

H

H

H

CH C

H

H H

C S H

H

H

Hydrogen BondsCompare the boiling points of H2O and H2S, shown in Table 2. These mol-ecules have similar sizes and shapes. However, the boiling point of H2O ismuch higher than that of H2S. A similar comparison of NH3 with PH3 canbe made.The greater the polarity of a molecule, the higher the boiling pointis. However, when hydrogen atoms are bonded to very electronegativeatoms, the effect is even more noticeable.

Compare the boiling points and electronegativity differences of thehydrogen halides, shown in Table 3. As the electronegativity differenceincreases, the boiling point increases.The boiling points increase somewhatfrom HCl to HBr to HI but increase a lot more for HF. What accounts forthis jump? The answer has to do with a special form of dipole-dipoleforces, called a hydrogen bond.


H

CH HC

H

H H

H

CC

H

H H

Table 3 Boiling Points of the Hydrogen Halides

Substance HF HCl HBr HI

Boiling point (°C) 20 −85 −67 −35

Electronegativity difference 1.8 1.0 0.8 0.5

Table 2 Comparing Dipole-Dipole Forces

hydrogen bond

the intermolecular force occurring when a hydrogen atom that is bonded to a highlyelectronegative atom of one molecule is attracted to twounshared electrons of anothermolecule

www.scilinks.orgTopic: Hydrogen BondingSciLinks code: HW4069

HHO

HHS

H

HH N

H

HH P


388 Chapter 11

Hydrogen Bonds Form with Electronegative AtomsStrong hydrogen bonds can form with a hydrogen atom that is covalentlybonded to very electronegative atoms in the upper-right part of the peri-odic table: nitrogen, oxygen, and fluorine. When a hydrogen atom bondsto an atom of N, O, or F, the hydrogen atom has a large, partially positivecharge. The partially positive hydrogen atom of polar molecules can beattracted to the unshared pairs of electrons of neighboring molecules. Forexample, the hydrogen bonds shown in Figure 10 result from the attrac-tion of the hydrogen atoms in the HIN and HIO bonds of one DNAstrand to the unshared pairs of electrons in the complementary DNAstrand. These hydrogen bonds hold together the complementary strandsof DNA, which contain the body’s genetic information.

Hydrogen Bonds Are Strong Dipole-Dipole ForcesIt is not just electronegativity difference that accounts for the strength ofhydrogen bonds. One reason that hydrogen bonds are such strong dipole-dipole forces is because the hydrogen atom is small and has only one elec-tron. When that electron is pulled away by a highly electronegative atom,there are no more electrons under it. Thus, the single proton of the hydro-gen nucleus is partially exposed.As a result, hydrogen’s proton is stronglyattracted to the unbonded pair of electrons of other molecules. The com-bination of the large electronegativity difference (high polarity) andhydrogen’s small size accounts for the strength of the hydrogen bond.

H

H

HH

H

H

H

HH

H

H

H

H

HCH3

O O

O

O

NN

NN

N

NN

N

N

N

NN N

N

N

Hydrogen bond

Figure 10Hydrogen bonding between basepairs on adjacent molecules of DNAholds the two strands together.Yet the force is not so strong thatthe strands cannot be separated.


Hydrogen Bonding Explains Water’s Unique PropertiesThe energy of hydrogen bonds is lower than that of normal chemicalbonds but can be stronger than that of other intermolecular forces.Hydrogen bonding can account for many properties. Figure 11 shows anexample of hydrogen bonding that involves oxygen. Water has uniqueproperties. These unique properties are the result of hydrogen bonding.

Water is different from most other covalent compounds because ofhow much it can participate in strong hydrogen bonding. In water, twohydrogen atoms are bonded to oxygen by polar covalent bonds. Eachhydrogen atom can form hydrogen bonds with neighboring molecules.Because of the water molecule’s ability to form multiple hydrogen bondsat once, the intermolecular forces in water are strong.

Another different characteristic of water results from hydrogen bond-ing and the shape of a water molecule. Unlike most solids, which aredenser than their liquids, solid water is less dense than liquid water andfloats in liquid water. The angle between the two H atoms in water is104.5°. This angle is very close to the tetrahedral angle of 109.5°. Whenwater forms solid ice, the angle in the molecules causes the special geom-etry of molecules in the crystal shown in Figure 11. Ice crystals have largeamounts of open space, which causes ice to have a low density.

The unusual density difference between liquid and solid waterexplains many important phenomena in the natural world. For example,because ice floats on water, ponds freeze from the top down and not fromthe bottom up. Thus, fish can survive the winter in water under an insu-lating layer of ice. Because water expands when it freezes, water seepinginto the cracks of rock or concrete can cause considerable damage due tofracturing. You should never freeze water-containing foods in glass con-tainers, which may break when the water freezes and expands.


Figure 11Water molecules arepulled together by fairly strong hydrogenbonds, which result in the open crystal structureof ice.

Solid Water

Liquid Water


Bonding of Water• In the diagram of 4 water molecules below, notice

the hydrogen ‘bonds’ labeled. Where are the polar covalent in this picture?


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Why do the water molecules line up

this way? !

Hint: You REALLY need to know how

to draw this diagram!!

Important Graph

• The graph below shows boiling points of Group 15, 16, and 17 elements bonded with hydrogen. Look at it carefully ... maybe even draw it in your notes ...


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• Why are the H2O, HF, and NH3 boiling points so much higher than expected?

summary

• In order of increasing strength of force:

• van der Waal (nonpolar molecules)

• dipole-dipole (polar asymmetric molecules)

• Hydrogen ‘bond’ (H: F, O, N)

• IonicStronger forces of attraction

Molecular Math

• Warm-up Exercise!

• What percent of this class is girls?

• What percent of this class is boys?

• What do the percent of boys and percent of girls have to add up to?

• What is the equation for calculating percent?

Topic 6

Calculating Percent Composition

Example 1:

• Suppose a sample of a compound contains 7.85 g Nitrogen and 17.92 g Oxygen. What is the percent composition of each element in this compound? (Use Table T)

Calculating percent composition

Example 2:

• A compound contains 2 elements, sulfur and oxygen. In a sample of the compound, the mass of sulfur is 6.41 grams. The mass of oxygen is 9.60 grams. What is the percent composition of the elements in this compound?

Determining Empirical Formula

• We’ve talked about this before, but you not remember, so here’s a refresher ...

The Empirical Formula of a compound shows the relative number of atoms of each element in the Simplest Whole Number Ratio

!For example, the molecular formula C6H12O6 can be “reduced” to the

empirical formula C1H2O1 !

The molecular formula C6H15 can be “reduced” to the empirical formula __________________ .

!Ionic compounds are ALWAYS represented in their empirical form,

called a Formula Unit.

Empirical Examples• Now try a few yourself ... Write the empirical

formula for each of the following:

• C2H6

• C2H4

• C2H2

• NO2

• C3H9

• N3O

• C3PO

Finding empirical formula from percentage composition

• Unfortunately, this is not an equation on Table T. You must MEMORIZE the following steps in order to solve these problems.

• Let’s go through an example and then you’ll have a chance to try one for yourselves at your table groups.

Example #1

• You have a compound and 30.46% of it is Nitrogen and 69.54% of it is Oxygen. To determine its empirical formula:

• Assume you have a 100 gram sample of the compound. That means you have 30.46 grams of Nitrogen and 69.54 grams of Oxygen.

• Convert grams to moles for both N and O.

• Convert the # of moles of N and O from step 2 into the simplest whole number ratio. (Divide both # of moles by whichever one is the smaller number.)

• This simple ratio becomes the subscripts in the empirical formula.

Example #2

• Try this one at your table groups.

• A compound contains 0.467 grams of Sulfur and 1.033 grams of Chlorine. What is the empirical formula of the compound?

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• Remember to follow the steps!

How to find Molecular Formula

• If you know BOTH the empirical formula and the gram formula mass of a compound, you can find the molecular formula of the compound.

• Once again, let’s outline the procedure for an example and then try one at your table groups...

Example #1 Molecular Formula

• The empirical formula of a compound is NO2, and the molecular mass is 92.0 grams. Find the molecular formula.

• Determine the gram formula mass of the empirical formula (Periodic Table).

• Divide the molecular mass by the empirical formula mass, then round that answer to the nearest whole number.

• Multiply the subscripts in the empirical formula by that number to get the correct molecular formula.

Example #2

• A compound used to whiten teeth has the empirical formula of HO. The molecular mass of this compound is 34.02 amu. What is its molecular formula?

Unit Overview• Topic 1 - Bonding (Ionic, Covalent, and Metallic)

• Topic 2 - Organic Chemistry (naming, and drawing structures of alkanes, alkenes, and alkynes).

• Topic 3 - Polar & Non-Polar Bonds (electronegativity affects polarity of a bond. The > the difference in electronegativity, the more polar the bond.)

• Topic 4 - Polar & Non-Polar Molecules (symmetry, non-bonded pairs of electrons)

• Topic 5 - Intermolecular Forces (van der Waal’s [London Dispersion / Dipole-Dipole], Hydrogen ‘Bond’, Ionic)

• Topic 6 - Molecular Math (Percent Composition, Empirical Formula, Molecular Formula)

unit 8 - bonding organic - dr. g's...

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