chemistry 20

Chemistry 20

PowerPoint presentation byR. Schultz

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Chapter 2

2.1 Three-Dimensional Structures

• Ionic• Recall electron transfer (Chapter 1)

• As ions form, they group – not in pairs – in stable crystals, of anions and cations in an arrangement so that each cation is surrounded by anions and the reverse

• Crystal is called (ionic) crystal lattice

Na• + Cl •••

••

•• •

••

••

••Cl•Na + ‾

electron transfer

+


• NaCl crystal lattice:

• In the crystal there are equal numbers of Na+ and Cl‾

fig 2.2, page 48

fig 2.1 A, page 48

photo of NaCl crystal: note its cubic shape – matches theoretical model


• Formula is written NaCl – this is the formula unit – the simplest ratio of anions to cations in the lattice – same as formula from balancing charges

• Different ionic compounds have different crystal shapes, 2 factors:– Relative ion size (see fig 2.4A, page 69)– Relative ion charge (see fig 2.4B, page 69)


• There are no molecules of ionic compounds, the term is formula unit


• Molecular• The smallest unit of a molecular compound is

a molecule• Molecular structures can be determined using

Lewis Diagrams along with the VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) Theory

• Theory based on repulsion between valence electron pairs, both lone (LP) and bonding (BP) pairs


• Valence electron pairs move as far apart as possible to minimize repulsion

• Repulsion order:

LP-LP > LP-BP > BP-BP• Shapes around central atoms can be

determined using the following chart (an improvement of chart page 55 of your text)

2.1 Three-Dimensional StructuresVSEPR Theory – Summary Chart

VSEPR Class

Name of molecular

shape

Types of electron

pairs

Example Lewis dot diagram VSEPR Diagram

Model

AX2

linear

all BP

0 LP 2 BP

CO2

AX3

trigonal planar

all BP

0 LP 3 BP

CH2O

AX2E

trigonal planar (bent)

1 LP, 2 BP

SO2

AX4

tetrahedral

all BP

0 LP 4 BP

CH4

AX3E

tetrahedral (trigonal

pyramidal)

1 LP, 3 BP

NH3

AX2E2

tetrahedral (bent)

2 LP, 2 BP

H2O

O C O

•• ••

°°

°°

°° °°

°°

°°

H C

H •x

•x

•• O °°

°°

°° •• O

°°

°°

°° S •• •• O °°

°°

°°

C •x •x

•x

•x H

H

H

H

N ••

•x

•x

•x

H

H H

O °° °°

° • ° •

H

H

not necessary to state


• Work through examples using chart as reference:

CH4, NH3, H2O, C2H4, SO2, C2H2

• VSEPR Worksheet (BLM 2.1.8B) – “prelab”• Models tomorrow


H C H••

••

••

••

H

H

Each electron pair around C is a BP.

They move to get as far apart as possible.

Shape: tetrahedral

Bond angle: 109º

flat on page

out of page

behind page

0 LP’s,4 BP’s

whenever there are 0 LP’s and 4 BP’s around central atom shape will be tetrahedral


N••

••

•• ••HHH

1 LP, 3 BP’s around Nall repel each other

shape like tetrahedral, but one bond space taken up by LP

shape name: pyramidal (book calls it “tetrahedral (trigonal pyramidal)”)

bond angle approx 107º

a little smaller than 109º since LP-BP repulsion greater than BP-BP

whenever there is 1 LP and 3 BP’s around central atom, shape will be pyramidal


O••

•• ••••H

H2 LP’s, 2 BP’s around Oall repel each other

shape like tetrahedral, but 2 bond spaces taken up by LP’s

shape name: bent or v-shaped (book calls it “tetrahedral (bent)”)


a little smaller than 109º since LP-BP repulsion greater than BP-BP – there are now 2 LP-BP sets

OHH

whenever there are 2 LP’s and 2 BP’s around central atom, shape will be v-shaped or bent


C C•• ••

••

•• ••

••H

HH

H

New rule: count multiple bonding pairs as if they were single bonding pairs

C = C

H

H

H

H

2 central atoms, each with 0 LP’s, 3 BP’s around themall repel each othershape name: trigonal planarbond angle 120º

flat on page!

120º

120º120º

120º

120º

120º

whenever there are 0 LP’s and 3 BP’s around central atom, shape will be trigonal planar


S••

••

••OO ••••

•• •••• ••

1 LP, 2 BP’s around Sall repel each other

shape like trigonal planar, but 1 bond space taken up by LP

shape name: trigonal planar (bent)


whenever there is 1 LP and 2 BP’s around central atom, shape will be trigonal planar (bent)


C C•• •••• HH •••• 0 LP’s, 2 BP’s around each Call repel each othershape name: linearbond angle approx 180º

CCH H

whenever there is 0 LP’s and 2 BP’s around central atom, shape will be linear


• You must know the chart – it’s not just memory

• know 3 lines; the rest can be worked out from that

• Do worksheet BLM 2.1.8C


• Polar Molecules• Chapter 1 – polar covalent bonds• Follow-up → polar molecules

• Whether or not a molecule is polar affects its bonding and therefore its chemical and physical properties


• To be polar, molecule must have at least 1 polar covalent bond

• Having 1 or more polar covalent bonds doesn’t necessarily make molecule polar

• Direction matters

OHH

O = C = O

polar non-polar


e.g. CO2

e.g. CH2O

e.g. CH3Cl

e.g. CH4

e.g. NH3

e.g. H2O

e.g. COS

e.g. SO3

Note: arrow directions are relative!


• Polarity of more complex molecules can be determined using the same methods

• e.g. C2H5F

• Do worksheet BLM 2.1.12B

C C

H

H

F

If F had been an H, dipole would’ve been in opposite direction and net dipoles from the 2 C’s in opposite directions, making it non-polar, but ………


• Generalizations for later:– all hydrocarbons & halocarbons, CxHy, e.g. CH4,

C3H8, C2F6 will be non-polar

– all monosubstituted hydrocarbons & halocarbons CxHyX, e.g. C2H5Br, CH3OH, CHI3 will be polar

• These will save you a lot of time later in chapter


• Network Solids (also called Covalent Networks) covalent bonds – strongest chemical bonds

• atoms bonded together by covalent bonds in a molecule made up of gigantic # of atoms

• Examples:• Diamond, C(s) – Each C is covalently bonded

to 4 other C’s in a arrangementtetrahedral

Fig 2.17B, page 601 diamond is 1 gigantic molecule!


• Graphite, C(s) – Each C is covalently bonded to 3 other C’s in a arrangement forming 2-dimensional layers

• Weaker bonds between layers

• Why is diamond hard and brittle with a very high melting point (mp) while graphite is slippery with a somewhat lower mp?

trigonal planar

fig 2.17A, page 60Each layer of crystal is 1 molecule


• other allotropes of C(s): fullerenes, carbon nanotubes, graphene (Read pages 60, 61)

• Silicon carbide, SiC(s) – structure like graphite, but mixture of Si and Chigh mp, hard, brittle like diamond – used in grindstones, carbide drill bits and saw blades

• Silicon dioxide, SiO2(s) – main ingredient of sand – high mp, hard, brittle

fig 2.19, page 61

2.2 Intermolecular Forces

• Intramolecular Forces (intramolecular bonds) – bonds between atoms in a molecule or between atoms in a polyatomic ion, a.k.a. :

• Intermolecular forces (intermolecular bonds) – bonds among molecules in molecular substances

covalent bonds


OH

H

OH

Hintermolecular forces

intramolecular forces

(polar) covalent bonds

intermolecular forces are important in understanding the differing boiling points (bp’s) and mp’s of molecular compounds


• Types of intermolecular forces:– dipole-dipole forces– hydrogen bonding– London (dispersion) forces

• dipole-dipole forces (polar molecular compounds only)

Also called van der Waals Forces


• Recall polar molecules, e.g. H2O

OHH molecular dipole

δ-

δ+

bond dipoles

O

HH

O

HH

O H

H

OH

H

OH

H

O

HH

OH

HO

HH

O H

H

OH

H

OH

H

O H

H

fig 2.21 page 63

“+” pole of 1 molecule attracts “–” poles of neighbours (much weaker than ionic bonds)


• hydrogen bondinga special kind of dipole-dipole force that occurs when H is covalently bonded to a very electronegative atom (F, O, or N)

• easy way to remember – instead of hydrogen bonding, think hydrogen FONding

• strongest intermolecular bond in most situations

• H has only 1 electron; when bonded to a very electronegative atom that electron is far


from the H, making it very +’ve.

• The attraction of the +’ve H for electron pairs on the other atom is a hydrogen bond

H O•• ••••

••H

H O•• ••

••

••

H

very +’ve

H O•• ••

••

••

H

hydrogen bond

remember that actual shape around O is v-shaped leading to …


Fig 2.23A, page 65

This accounts for the hexagonal configuration of snowflakes, for the fact that ice floats on water, and for water’s high surface tension

Remember the pennies in the glass?


• London (Dispersion) Forces• Many molecular compounds have neither

dipole-dipole forces nor hydrogen bonding – they are non-polar and there is no charge-charge interaction, therefore…….

• According to theory there would be no attraction between these molecules and their only possible state would be gas (at any temperature)


• London Dispersion Forces (LDF) are present in ALL molecular compounds

• some molecular compounds have LDF along with the other forces

• illustration: charged balloon will be attracted to the uncharged wall because charges are induced in the wall as shown

• motion of electrons in atom canproduce temporary dipoles

wall

fig 2.25, page 68


• the temporary dipoles induce temporary dipoles in neighbouring molecules

• attraction between these temporary dipoles and the induced dipoles is the source of London Dispersion Forces

• Factors affecting strength of LDF:– total number of electrons in the molecule– shape of the molecule*

and their neighbouring molecules


• Read about the importance of hydrogen bonding in the DNA double helix and in myoglobin in muscle tissue

• page 69-70

2.3 Relating Structures and Properties

• State of Matter (mp and bp)

page 72


• raising temperature is raising Ek of particles

• the stronger the appropriate bonding, the higher the mp and bp of the substance

• appropriate bonding:network covalents – covalent bondingionics – ionic bondingmetals – metallic bondingmoleculars – LDF, dipole-dipole, H-bonding(note: even though moleculars have covalent bonding this has nothing to do with mp and bp)


• Relative bond strengths:

covalent > ionic > metallic > hydrogen > dipole-dipole > LDF (LDF varies

greatly)

strongest weakest


chart, page 72

high mp & bp; metallic bonding is strong – no differentiation

high mp & bp; ionic bonding is strong – no differentiation

low mp & bp; intermolecular forces are weak - differentiation


• Hydrogen bonding has a great affect on boiling points:

10 e‾ 18 e‾ 36 e‾ 54 e‾

effect of H bonding

Why do bp’s increase as you go from left to right?

How come group 14 always is lowest?

Where water would be expected to be without H bonding

fig 2.33, page 74


• Examples:• Predict which will have the higher boiling point

in each pair:

• a) C2H6 or C3H8 both molecular: intermolecular forces only

LDF 18 e‾ 26 e‾DDF no noH bonds no no

C3H8 will have higher bp due to stronger LDF


• b) CH3F and CH3OH LDF 18 e‾ 18 e‾DDF yes yesH bonds no yes

Both molecular: intermolecular forces only

CH3OH will have the higher bp because of hydrogen bonding

Do worksheet BLM 2.3.3B


• Extension: Rank the following substances in order of bp. State relevant bond types for each

• CH3OH, SiC, PH3, LiF, CH3F, Mg

Step 1: Identify moleculars – they will be lowest

M M M

LDF 18 e‾ no 18 e‾ no 18 e‾ noDDF yes no no no yes noH bonds yes no no no no no

Other covalent ionic metallic (network)

Answer:

PH3, CH3F, CH3OH, Mg, LiF, SiC


• Note: on your test the substances will always be in pairs as on your previous worksheet

• unlike the previous worksheet, not everything will necessarily be molecular – i.e. you may get a molecular with an ionic, an ionic with a covalent network, 2 moleculars; any grouping is possible


• Why are metals malleable according to the metallic bonding model? (possible question on next test)

• Malleable means workable – bends without breaking; dents without shattering

• metallic bonding model: an array of cations in a sea of valence electrons

• because bonds are between cations and sea of electrons rather than between cations and cations, cations may be moved in the sea of electrons without breaking bonds


• Ionic compounds are not malleable – they shear along lines as diagram indicates

fig 2.36, page 75


• Why do metals conduct electricity according to the metallic bonding model? (possible question on next test)

+e-

e-

e-

e-

magnified section of wire

e-e-

+


• electricity is the flow of electrons• metallic bonding model: an array of cations in

a sea of valence electrons• since valence electrons are loosely held to

individual nuclei, they are delocalized – as a new electron enters the wire, another electron must leave to keep wire neutral


• Conductivity (others)

• solid ionic:

page 76


• aqueous ionic:

page 77


• molecular compounds:

• No – not even if polar molecular

page 77


• Investigation 2.E, page 78


• Chapter Review: page 82, questions 1 - 5,7, 9 - 12, 13a, b, d, e, i, k, n, p, t, 14b, c, d, f, h, 15

• Work on these and be prepared to discuss in class


• T

chemistry 20

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