Intermolecular forces- IMF

Intermolecular forces- IMF
12/10/99 Intermolecular forces- IMF Covalent bond – chemical bond (strong) Intermolecular attraction - physical bond (weak forces) Application of Core principles Of chemistry Edexcel new Specification

Content of the specification Intermolecular forces A
12/10/99 A Demonstrate an understanding of the nature of intermolecular forces resulting from interactions between permanent dipoles, instantaneous dipoles and induced dipoles (London forces) and from the formation of hydrogen bonds. B Relate the physical properties of materials to the types of intermolecular force present, eg: i The trends in boiling and melting temperatures of alkanes with increasing chain length. ii The effect of branching in the carbon chain on the boiling and melting temperatures of alkanes. iii The relatively low volatility (higher boiling temperatures) of alcohols compared to alkanes with a similar number of electrons. iv The trends in boiling temperatures of the hydrogen halides HF to HI.

C Carry out experiments to study the solubility of simple molecules in different solvents D Interpret given information about solvents and solubility to explain the choice of solvents in given contexts, discussing the factors that determine the solubility including: i The solubility of ionic compounds in water in terms of the hydration of the ions ii The water solubility of simple alcohols in terms of hydrogen bonding iii The insolubility of compounds that cannot form hydrogen bonds with water molecules, eg polar molecules such as halogenoalkanes iv The solubility in non-aqueous solvents of compounds which have similar intermolecular forces to those in the solvent.

Intermolecular Forces
IM forces originate from interactions between charges, partial charges, and temporary charges on molecules. IM forces are relatively weak because of smaller charges and the distance between molecules.

Dipole forces Polar covalent molecules are sometimes described as "dipoles", meaning that the molecule has two "poles". One end (pole) of the molecule has a partial positive charge while the other end has a partial negative charge. The molecules will orientate themselves so that the opposite charges attract principle operates effectively.

Contains polar molecules Liquid shows deflection Contains non-polar molecules Liquid shows no deflection

Deflection of a polar liquid (water) under the influence of a charged rod.
Note that symmetrcal molecules like tetracloromethane and Cyclohexane do not deflect as they are not polar although they have polar bonds deflection of water

Intermolecular forces- IMF
Introduction The physical properties of melting temperature, boiling temperature, vapor pressure, evaporation, viscosity, surface tension, and solubility are related to the strength of attractive forces between molecules. These attractive forces are called Intermolecular Forces. The amount of "stick togetherness" is important in the interpretation of the various properties. There are four types of intermolecular forces. Most of the intermolecular forces are identical to bonding between atoms in a single molecule. Intermolecular forces just extend the thinking to forces between molecules and follows the patterns already set by the bonding within molecules.

electrostatic attraction between dipoles, i
electrostatic attraction between dipoles, i.e the attraction between the +ve end of the molecule and –ve end of another molecule

Polar Molecule If the difference in electronegativity is not so great, however, there will be some degree of sharing of the electrons between the two atoms. The result is the same whether two ions come together or two atoms come together:

The Relationship Between Electronegativity and Bond Type

Intermolecular forces
1.Van Der Waal’s Forces (London forces) 2.Hydrogen bonding 3.Ion-dipole forces Van der Waals’ forces Dipole-Dipole Interaction Dipole- Induced Dipole Interaction Instantaneous Dipole- Induced Dipole Interaction

Intermolecular forces
12/10/99 Intermolecular forces Intermolecular forces ( IMF) are attractive forces between molecules that occur when there is a variation in the electron distribution in a molecule. Intermolecular forces are weaker than the weakest covalent bonds. Intermolecular forces arise when a partially negative charge on a molecule is attracted to a partially positive charge on another molecule.

A. Definition of IMF Attractive forces between molecules.
Much weaker than chemical bonds within molecules.

12/10/99 Intermolecular Forces The forces holding solids and liquids together are called intermolecular forces. Intermolecular forces are much weaker than ionic or covalent bonds. Example: 16 kJ/mol to vaporize HCl compared to 431 kJ/mol to break HCl into its elements. When a substance melts or boils, the intermolecular forces are broken (not the covalent bonds).

12/10/99 Intermolecular Forces The orientation of polar molecules in an electric field Boiling temperature reflects intermolecular force strength. A high boiling temperature indicates strong attractive forces. A high melting temperature also reflects strong attractive forces. Electric field ON Electric field OFF

Polarity and Boiling Point
The polarity of the molecules determines the forces of attraction between the molecules in the liquid state. Polar molecules are attracted by the opposite charge effect (the positive end of one molecule is attracted to the negative end of another molecule. Molecules have different degrees of polarity as determined by the functional group present. The greater the forces of attraction the higher the boiling temperature or the greater the polarity the higher the boiling temperature.

Types of Intermolecular Forces
In pure substances: London forces (dispersion forces) -- very weak “instantaneous induced dipole” forces between molecules H-bonding -- especially strong dipole-dipole forces for compounds with H-F, H-O, or H-N bonds dipole-dipole forces -- between polar molecules (e.g. SO2, PF3) Forces within mixtures (in addition to the above): ion-dipole -- between ionic and polar substances ion-induced dipole -- between ionic and non-polar substances dipole-induced dipole -- between polar and non-polar substances

Prof. Fritz London Johannes Diderik van der Waals

Types of Van der Waals Forces
London forces (dispersion forces) very weak “instantaneous induced dipole” forces between molecules Van der Waals Forces Types of Van der Waals Forces 1) Dispersion 2) Dipole – Dipole Interaction 2) Dipole Induced Dipole Interaction 3) Instantaneous Dipole- Induced Dipole Interaction

Interacting Nonpolar Molecules
Dispersion forces (London dispersion forces) are intermolecular forces caused by the presence of temporary dipoles in molecules. A instantaneous dipole (or induced dipole) is a separation of charge produced in an atom or molecule by a momentary uneven distribution of electrons.

Dispersion Forces (Instantaneous dipoles)
Cl-Cl e- e- Cl-Cl e- e- d- d+ e- e- d- d+ e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- INDUCED DIPOLE TEMPORARY DIPOLE non-polar non-polar Dispersion (weakest and very short-lived)

Illustrations A instantaneous dipole (or induced dipole) is a separation of charge produced in an atom or molecule by a momentary uneven distribution of electrons.

Prof. Fritz London London Dispersion Forces –
significant only when molecules are close to each other Due to electron repulsion, a temporary dipole on one atom can induce a similar dipole on a neighboring atom Prof. Fritz London

London Forces (Dispersion)
Induced dipoles (Instantaneous ) Strength is surface area dependent More significant in larger molecules All molecules show dispersion forces Larger molecules are more polarizable

The ease with which an external electric field can induce a dipole (alter the electron distribution) with a molecule is referred to as the "polarizability" of that molecule The greater the polarizability of a molecule the easier it is to induce a momentary dipole and the stronger the dispersion forces Larger molecules tend to have greater polarizability

Their electrons are further away from the nucleus (any asymmetric distribution produces a larger dipole due to larger charge separation) The number of electrons is greater (higher probability of asymmetric distribution) thus, dispersion forces tend to increase with increasing molecular mass Dispersion forces are also present between polar/non-polar and polar/polar molecules (i.e. between all molecules)

Strength of Dispersion Forces
The strength of dispersion forces depends on the polarizability of the atoms or molecules involved. Polarizability is a term that describes the relative ease with which an electron cloud is distorted by an external charge. Larger atoms or molecules are generally more polarizable than small atoms or molecules.

London Dispersion Forces

Instantaneous and Induced Dipoles
2,2 – dimethylprppane Boiling temperature 9.5 C pentane Boiling temperature C

Molar Mass and Boiling Temperature
Relative Molecular Mass and Boiling Temperature of Common Species. Halogen No of Lone pairs M Tb(K) Noble Gas He 1 2 4 F2 3 38 85 Ne 20 27 Cl2 71 239 Ar 40 87 Br2 160 332 Kr 84 120 I2 254 457 Xe 131 165 Rn 211 r r

Ar Example e- H Cl d+ d- e- e- e- e- e- d+ d- e- e- e- e- e- e- e- e-
non-polar A DIPOLE (it’s polar) INDUCED DIPOLE Dipole – Induced Dipole (weak and short-lived)

The Effect of Shape on Forces
boiling temperature 28 C boiling temperature 9 C boiling temperature 37 C

Practice Rank the following compound in order of increasing
boiling temperature CH3OH, CH3CH2CH2CH3, and CH3CH2OCH3

Practice Rank the following compound in order of increasing boiling temperature. CH3OH, CH3CH2CH2CH3, and CH3CH2OCH3 MM IM Forces CH3OH 32.0 London and H-bonding CH3CH2CH2CH3 58.0 London, only CH3CH2OCH3 60.0 London and Dipole-dipole

Dipole-dipole interactions

Permanent Dipole-Permanent Dipole forces

Polar molecules can interact with ions:
Ion - Dipole Interactions

Ion-Dipole Forces Interaction between an ion and a dipole.
12/10/99 Ion-Dipole Forces Interaction between an ion and a dipole. Strongest of all intermolecular forces. Example: Na+ and Cl- ions dissolved in water. Ion-dipole attractions become stronger as either the charge on the ion increases, or as the magnitude of the dipole of the polar molecule increases.

Weaker Intermolecular Forces
Ion-Dipole Forces An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a neutral molecule that has a dipole. Most commonly found in solutions. Especially important for solutions of ionic compounds in polar liquids. A positive ion (cation) attracts the partially negative end of a neutral polar molecule. A negative ion (anion) attracts the partially positive end of a neutral polar molecule

Dipole-Dipole Interactions
Found in PC molecules. Stronger than LDFs LDFs and Dipole-Dipole are also called as Van Der Waals forces Ex: ICl

Dipole – Dipole attractions
12/10/99 Polar covalent molecules are sometimes described as "dipoles", meaning that the molecule has two "poles". One end (pole) of the molecule has a partial positive charge while the other end has a partial negative charge. The molecules will orientate themselves so that the opposite charges attract principle operates effectively.

between polar molecules
Dipole-Dipole Forces between polar molecules Exist between neutral polar molecules. Weaker than ion-dipole forces. If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity.

Dipole-Dipole Interactions
Dipole-dipole interactions are attractive forces between polar molecules. An example is the interaction between water molecules. The hydrogen bond is a special class of dipole-dipole interactions due to its strength.

Dipole-Dipole Forces + -

Effect of Dipole-Dipole Force
Polar molecules have dispersion forces and dipole-dipole forces. Effects can be seen in boiling and melting temperature. Tb C Tm C

Dipole-Dipole Force Occurs in polar molecules which have permanent dipoles, so attraction between molecules is always present.

Dipole-Dipole Forces H Cl H Cl H Cl Dipole-dipole (Polar molecules) δ+
Alignment of polar molecules to two electrodes charged + and δ– Forces compared to ionic/covalent are about 1 in strength compared to a scale of 100, thus 1% δ+ δ– δ+ δ– δ+ δ– H Cl H Cl H Cl

Dipole Dipole Interactions

Permanent Dipole A permanent dipole exists in all polar molecules as a result of the difference in the electronegativity of bonded atoms.

12/10/99 Dipole-Dipole Forces Molecular masses, Dipole moments and Boiling points of several simple organic molecules Molecule Molecular mass Dipole moment: D Boiling temperature CH3CH2CH3 44 0.1 231 CH3OCH3 46 1.3 248 CH3Cl 50 1.9 249 CH3CHO 2.7 294 CH3CN 41 3.9 355

Weakest of all intermolecular forces… London dispersion forces exist between all molecules! How it is formed ?? Temporary asymmetrical distribution/ random arrangement of electrons/ charge(density) OR instantaneous/temporary dipole (these produce) induced dipoles

Dispersion Force Dispersion force (London force) is present in all molecules and atoms and results from changes in electron- locations.

Instantaneous Dipoles
Charge separation in one creates charge separation in the neighbors.

Dispersion Force Strength
Polarizability indicates how readily an electron cloud can be distorted. The larger the atom, the more loosely it holds the electrons in its outermost shell, and the more they can be distorted. The more polarizable the atom, the stronger are the van der Waals interactions Large atoms with large electron clouds tend to have stronger dispersion forces. Large molecules tend to have stronger dispersion forces.

12/10/99 London Dispersion Forces “instantaneous dipoles” The larger the molecule (the greater the number of electrons) the more polarizable or the easier it is to create instantaneous dipoles. London dispersion forces increase as molecular weight increases.

Boiling Temperature of The Halogens and the Noble Gases Halogen R.M.M Tb K Noble Gas R.A.M F2 Cl2 Br2 I2 38 71 160 254 85.1 238.6 332 457.6 He Ne Ar Kr Xe 4 20 40 84 131 4.6 27.3 87.5 121 166

Polarizability is a measure of how the electron cloud around an atom responds to changes in its electronic environment.

12/10/99 London Dispersion Forces London dispersion forces depend on the shape of the molecule. The greater the surface area available for contact, the greater the dispersion forces. London dispersion forces depend on number of electrons, The more electrons in a molecule, the grater London dispersion forces London dispersion forces between branched nonpolar molecules are lower than the forces between long nonpolar molecules.  More branching in alkane decreases relative surface area so less contact between (neighbouring) molecules and reduces london forces

If you think of the unbranched alkane pentane as a cigar and branched pentane as a tennis ball, you can see that branching decreases the area of contact between molecules: Two cigars make contact over a greater area than do two tennis balls. Thus, if two alkanes have the same molecular weight, the more highly branched alkane will have a lower boiling point. As they highly brached alkanes will have less contact with neighboring molecules

Tb(°C) Tm(°C) Number of carbon atom(s) Straight-chain alkane
Physical Properties of Alkanes Number of carbon atom(s) Straight-chain alkane Tb(°C) Tm(°C) Density at 20°C (g cm–3) 1 2 3 4 5 6 7 8 9 10 Methane Ethane Propane Butane Pentane Hexane Heptane Octane Nonane Decane –161 –89 –42 36 69 98 126 151 174 –183 –172 –188 –135 –130 –95 –91 –57 –54 –30 — 0.626 0.657 0.684 0.703 0.718 0.730 At R.T., C1 – C4: gases ; C5 – C17: liquids ; > C18: waxy solid

Intermolecular Forces and Changes of State
This table shows the boiling temperature of some common straight-chain alkanes. As the number of carbon atoms increases, the boiling temperature increases. Boiling temperature of common straight- chain alkanes Tb C

Butane has a higher boiling temperature than
2-methylpropane. This is because butane has A stronger C–H bonds. B more electrons. C a larger surface area. D hydrogen bonds. An electric field can affect the direction of a stream of some liquids. Which of these liquids would be affected by an electric field? A 1-chloropropane B Pentane C Tetrachloromethane D Cyclopentane

Which of these isomers has the highest boiling temperature?
A B C D

Strength of Van der Waals’ Forces
Type of interaction Magnitude (kJ mol-1) Dipole-dipole 5-25 Dipole-induced dipole 2-10 Instantaneous dipole-induced dipole

Hydrogen bonding Hydrogen bonding is usually stronger than normal dipole forces between molecules. Of course hydrogen bonding is not nearly as strong as normal covalent bonds within a molecule - it is only about 1/10 as strong. This is still strong enough to have many important ramifications on the properties of water.

London vs Hydrogen Bonding
Hydrocarbon Alcohol Molecular Formula Molar Mass Tb (oC) CH4 16.04 -161.5 CH3CH3 30.07 -88 CH3OH 32.04 64.5 CH3CH2CH3 44.09 -42 CH3CH2OH 46.07 78.5 CH3CH(CH)CH3 58.12 -11.7 CH3CH(OH)CH3 60.09 82 CH3CH2CH2CH3 -0.5 CH3CH2CH2OH 97

The difference in boiling temperature between methane (Tb = 109 K) and
ethane (Tb = 185 K) is best explained by the different numbers of A protons. B electrons. C covalent bonds. D hydrogen bonds.

Butane has a higher boiling temperature than 2-methylpropane
Butane has a higher boiling temperature than 2-methylpropane. This is because butane has A stronger C–H bonds. B more electrons. C a larger surface area. D hydrogen bonds.

Hydrogen “Bonding” This IM force is a misnomer since it’s not an actual bond. Occurs between molecules in which H is bonded to a highly electronegative element (N, O, F), leading to high partial positive and partial negative charges. It’s a “super” dipole-dipole force.

Effect of H “Bonding” Hydrogen “bonding” is a very strong intermolecular force. Molecules with H “bonding” have much higher than expected melting and boiling points.

The ethanol molecule contains a polar O—H bond.

H “Bonding” in Ethanol & Water

Physical properties of water related with IMF
For most substances, solids are more dense than liquids. This is not true for water. Water is less dense as a solid Ice floats on liquid water! Strong hydrogen bonds formed at freezing lock water molecules away from each other When ice melts, the structure collapses and molecules move closer together. This property plays an important role in lake and ocean ecosystems Floating ice often insulates and protects animals and plants living in the water below.

Hydrogen Bonds in ice and liquid water
In liquid water each molecule is hydrogen bonded to approximately 3.4 other water molecules. In ice each each molecule is hydrogen bonded to 4 other molecules. Compare the two structures below. Notice the empty spaces within the ice structure Water Ice

Hydrogen Bonds in liquid water
Hydrogen bonds are much weaker than covalent bonds. However, when a large number of hydrogen bonds act in unison they will make a strong contributory effect. This is the case in water.

Hydogen Bonds in liquid water
Liquid water has a partially ordered structure in which hydrogen bonds are constantly being formed and breaking up.

Evaporation of Liquid Water
This animation shows how water molecules are able to break the forces of attraction i.e. the hydrogen bonds to each other and escape as the gas molecule. This is what is happening inside the gas bubble as it is rising to the surface to break and release the water gas molecules.

Polarity and Boiling Temperature
The polarity of the molecules determines the forces of attraction between the molecules in the liquid state. Polar molecules are attracted by the opposite charge effect (the positive end of one molecule is attracted to the negative end of another molecule. Molecules have different degrees of polarity as determined by the functional group present. The greater the forces of attraction the higher the boiling point or the greater the polarity the higher the boiling temperature.

The origin of hydrogen bonding
The molecules which have this extra bonding are: Notice that in each of these molecules: * The hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen to acquire a significant amount of positive charge.

Each of the elements to which the hydrogen is attached is not only significantly negative, but also has at least one "active" lone pair. Lone pairs at the 2-level have the electrons contained in a relatively small volume of space which therefore has a high density of negative charge. Lone pairs at higher levels are more diffuse and not so attractive to positive things.

Summary of IM Forces Types of Intermolecular Forces

Mr °C CH4 16 -161 SiH4 32 -117 GeH4 77 -90 SnH4 123 -50 NH3 17 -33
Boiling temperature of hydrides Mr °C CH SiH GeH SnH NH PH AsH SbH Mr °C H2O H2S H2Se H2Te HF HCl HBr HI Group IV GROUP VI Group V GROUP VII The values of certain hydrides are not typical of the trend you would expect

Boiling temperature of group IV hydrides
Mr BOILING POINT / C° 100 -160 140 50 The boiling temperature of the hydrides increase with molecular mass. CH4 has the lowest boiling temperature as it is the smallest molecule. PbH4 GeH4 SiH4 Larger molecules have greater intermolecular forces and therefore higher boiling temperature CH4

Boiling temperature of group V hydrides
Mr BOILING POINT / C° 100 -160 140 50 NH3 has a higher boiling temperature than expected for its molecular mass. There must be an additional intermolecular force. NH3

Hydrogen Bonding in Water
Consider two water molecules coming close together. The + hydrogen is so strongly attracted to the lone pair that it is almost as if you were beginning to form a co-ordinate (dative covalent) bond. It doesn't go that far, but the attraction is significantly stronger than an ordinary dipole-dipole interaction. Hydrogen bonds have about a tenth of the strength of an average covalent bond, and are being constantly broken and reformed in liquid water.

Boiling temperature of group VI hydrides
Mr BOILING POINT / C° 100 -160 140 50 H2O H2O has a very much higher boiling temperature for its molecular mass. There must be an additional intermolecular force.

Boiling temperature of group VII hydrides
Mr BOILING POINT / C° 100 -160 140 50 HF has a higher boiling temperature than expected for its molecular mass. There must be an additional intermolecular force. HF

The higher than expected boiling
Boiling temperature of hydrides Mr BOILING POINT / C° 100 -160 140 50 H2O The higher than expected boiling temperature of NH3, H2O and HF are due to intermolecular Hydrogen bonding HF NH3 GROUP IV GROUP V GROUP VI GROUP VII

Hydrogen iodide has a higher boiling temperature than hydrogen bromide
Hydrogen iodide has a higher boiling temperature than hydrogen bromide. This is because A the H–I bond is stronger than the H–Br bond. B hydrogen iodide has stronger London forces than hydrogen bromide. C hydrogen iodide has a larger permanent dipole than hydrogen bromide. D hydrogen iodide forms hydrogen bonds but hydrogen bromide does not.

The hydrogen halides The hydrogen halides are colourless gases at room temperature, producing steamy fumes in moist air. Hydrogen fluoride has an abnormally high boiling point for the size of the molecule (293 K or 20°C), and could condense to a liquid on a cool day.

Hydrogen fluoride Hydrogen fluoride's boiling temperature is higher than you might expect because it forms hydrogen Fluorine is the most electronegative of all the elements and the bond between it and hydrogen is very polar. The hydrogen atom carries quite a lot of positive charge ( +); the fluorine is fairly negatively charged ( -).

Hydrogen fluoride In addition, each fluorine atom has 3 very active lone pairs of electrons. Fluorine's outer electrons are at the 2-level, and the lone pairs represent small highly charged regions of space. Hydrogen bonds form between the + hydrogen on one HF molecule and a lone pair on the fluorine of another one.

The other hydrogen halides
The other hydrogen halides don't form hydrogen bonds. The other halogens aren't as electronegative as fluorine, and so the bonds in HX are less polar. As well as that, their lone pairs are at higher energy levels. That makes the lone pairs bigger, and so they don't carry such an intensely concentrated negative charge for the hydrogens to be attracted to.

Water as a "perfect" example of hydrogen bonding
Notice that each water molecule can potentially form four hydrogen bonds with surrounding water molecules. There are exactly the right numbers of + hydrogens and lone pairs so that every one of them can be involved in hydrogen bonding.

Water as a "perfect" example of hydrogen bonding
This is why the boiling temperature of water is higher than that of ammonia or hydrogen fluoride. In the case of ammonia, the amount of hydrogen bonding is limited by the fact that each nitrogen only has one lone pair. In a group of ammonia molecules, there aren't enough lone pairs to go around to satisfy all the hydrogens. In hydrogen fluoride, the problem is a shortage of hydrogens. In water, there are exactly the right number of each. Water could be considered as the "perfect" hydrogen bonded system.

Hydrogen bonding Comparison of Bond Lengths:
The graphic on the right shows a cluster of water molecules in the liquid state. Water is a polar molecule, with the oxygen (red) being the negative area and the hydrogen (white) being the more positive area. Opposite charges attract. The bond lengths give some indication of the bond strength. A normal covalent bond is 0.96 Angstroms, while the hydrogen bond length is is 1.97 A.

Introduction to Organic Molecules and Functional Groups

Which of the following compounds has the highest boiling temperature?
A CH4 B CH3Cl C HCHO D CH3OH

Methanol dissolves in water mainly due to the formation of new
A hydrogen bonds. B dipole-dipole forces. C London forces. D covalent bonds. Consider the following organic liquids: A ethanal B ethanol C tetrachloromethane D trichloromethane Each liquid is run from a burette. Which liquid would not be deflected significantly by a charged rod? A B C D

For parts (a) and (b), use your knowledge of intermolecular forces to predict the compound with the highest boiling temperature. (a) A HF B H2O C NH3 D CH4 (b) A 1-iodobutane B 1-chlorobutane C 2-methyl-2-iodopropane D 2-methyl-2-chloropropane

Consider the following compounds, P, Q, R and S.
CH3CH2CH2CH3 Compound P Compound Q CH3CH2CH2CH2Br Compound R Compound S The boiling temperatures of compounds P, Q, R and S increase in the order A P Q R S B R S P Q C Q S P R D Q P S R

Which of the following compounds shows hydrogen bonding in the liquid state?
A Hydrogen bromide, HBr B Hydrogen sulfide, H2S C Silane, SiH4 D Ammonia, NH3 The ability of a liquid to flow is linked to the strength of its intermolecular forces. Suggest which of these liquids flows the slowest when poured. A Propane-1,2,3-triol B Propane-1,2-diol C Pentane D Butane

Which of the following has dipole-dipole interactions between its molecules, but no
hydrogen bonding? A Methane, CH4 B Methanol, CH3OH C Ammonia, NH3 D Hydrogen iodide, HI

Which list below shows the compounds in order of increasing boiling temperature?
A CH4, HCl, HF B HF, CH4, HCl C HCl, HF, CH4 D HF, HCl, CH4 Which of the following has the highest boiling temperature? A Pentane, CH3CH2CH2CH2CH3 B Hexane, CH3CH2CH2CH2CH2CH3 C 2-methylbutane, CH3CH(CH3)CH2CH3 D 2-methylpentane, CH3CH(CH3)CH2CH2CH3

Which intermolecular forces exist between molecules of ethoxyethane?
A Instantaneous dipole – induced dipole only B Permanent dipole – permanent dipole only C Instantaneous dipole – induced dipole and hydrogen bonds D Instantaneous dipole – induced dipole and permanent dipole – permanent dipole

Miracle of Hydrogen bonding

Six dancers

Hydrogen Bonding gives rise to even higher boiling points
an extension of dipole-dipole interaction gives rise to even higher boiling points bonds between H and the three most electronegative elements, F, O and N are extremely polar because of the small sizes of H, F, N and O the partial charges are concentrated in a small volume thus leading to a high charge density makes the intermolecular attraction greater and leads to even higher boilingtemperature

Hydrogen bonding in alcohols
An alcohol is an organic molecule containing an -O-H group. Any molecule which has a hydrogen atom attached directly to an oxygen or a nitrogen is capable of hydrogen bonding. Such molecules will always have higher boiling points than similarly sized molecules which don't have an -O-H or an -N-H group. The hydrogen bonding makes the molecules "stickier", and more heat is necessary to separate them. Ethanol, CH3CH2-O-H, and methoxymethane, CH3-O-CH3, both have the same molecular formula, C2H6O.

Formation of hydrogen bonds
in methanol

In methoxymethane, the lone pairs on the oxygen are still there, but the hydrogens aren't sufficiently + for hydrogen bonds to form. Except in some rather unusual cases, the hydrogen atom has to be attached directly to the very electronegative element for hydrogen bonding to occur.

The boiling temperature of ethanol and methoxymethane show the dramatic effect that the hydrogen bonding has on the stickiness of the ethanol molecules: ethanol (with hydrogen bonding) 78.5°C methoxymethane (without hydrogen bonding) -24.8°C The hydrogen bonding in the ethanol has lifted its boiling temperature about 100°C.

It is important to realise that hydrogen bonding exists in addition to van der Waals attractions. For example, all the following molecules contain the same number of electrons, and the first two are much the same length. The higher boiling temperature of the butan-1-ol is due to the additional hydrogen bonding. Boiling temp.

Comparing the two alcohols (containing -OH groups), both boiling points are high because of the additional hydrogen bonding due to the hydrogen attached directly to the oxygen - but they aren't the same.

The boiling point of the 2-methylpropan-1-ol isn't as high as the butan-1-ol because the branching in the molecule makes the van der Waals attractions less effective than in the longer butan-1-ol.

Ethanol and methoxymethane have the same number of electrons, and a similar length to the molecule. The van der Waals attractions (both dispersion forces and dipole-dipole attractions) in each will be much the same. However, ethanol has a hydrogen atom attached directly to an oxygen - and that oxygen still has exactly the same two lone pairs as in a water molecule. Hydrogen bonding can occur between ethanol molecules, although not as effectively as in water. The hydrogen bonding is limited by the fact that there is only one hydrogen in each ethanol molecule with sufficient + charge.

Which of the following compounds has highest boiling temperature?
A CH4 B CH3Cl C HCHO D CH3OH Which substance has the strongest London dispersion forces? A H2O B H2S C H2Se D H2Te

Hydrogen bonding in organic molecules containing nitrogen
Hydrogen bonding also occurs in organic molecules containing N-H groups - in the same sort of way that it occurs in ammonia. Examples range from simple molecules like CH3NH2 (methylamine) to large molecules like proteins and DNA. The two strands of the famous alpha-helix in DNA are held together by hydrogen bonds involving N-H groups.

Formation of hydrogen bonds in ammonia

Formation of hydrogen bonding in - HF
 Hydrogen fluoride has a much higher boiling point than one would expect for a molecule with a relative molecular mass of 20  Fluorine has the highest electronegativity of all and is a small atom so the bonding with hydrogen is extremely polar

HF molecule very +ve F being very electronegative
F atom being small enough to approach very close to the H atom in the neighbouring molecule

Hydrogen Bonding A special case of dipole-dipole forces.
12/10/99 Hydrogen Bonding A special case of dipole-dipole forces. By experiments, the boiling pts. of compounds with H-F, H-O, and H-N bonds are abnormally high. The intermolecular forces are therefore abnormally strong. H-bonding requires… H bonded to a small, highly electronegative element (most important for compounds of F, O, and N) 2) an unshared electron pair on a nearby small, highly electronegative ion or atom (usually F, O, or N on another molecule).

Which intermolecular forces exist between molecules of ethoxyethane?
A Instantaneous dipole – induced dipole only B Permanent dipole – permanent dipole only C Instantaneous dipole – induced dipole and hydrogen bonds D Instantaneous dipole – induced dipole and permanent dipole – permanent dipole

Hydrogen Bonding Examples
12/10/99 Hydrogen Bonding Examples abnormally high B.P. 1) polar molecules 2) nonpolar molecules 3)

The polar water molecule interacts strongly with the polar O—H bond in ethanol

Methanol dissolves in water mainly due to the formation of new
A hydrogen bonds. B dipole-dipole forces. C London forces. D covalent bonds.

Ethanol is soluble in water. The best explanation for this is
A ethanol molecules form hydrogen bonds with water molecules. B ethanol molecules form London (dispersion) forces with water molecules. C ethanol molecules form permanent dipole interactions with water molecules. D ethanol and water are miscible liquids.

The boiling temperatures of some hydrides are given below.
compound Boiling temperature ? K HF 293 HCl 188 HBr 206 HI 238 H2O 373 *(a) Explain, by comparing the forces involved, why HI has a higher boiling temperature than HBr.

Question Boiling points between HI and Hbr ?
More london forces in HI than in HBr Because HI has more electrons So it needs more energy to separate HI than Hbr Boiling points between HCl and HF ? HF has hydrogen bonds but HCl has london (dispersion) forces and weak dipole-dipole forces Hydrogen bonding in HF is stronger than london (dispersion) forces and weak dipole-dipole forces of HCL So more energy is required to separate HF molecule than HCl

Complication What is dipole moment ???
Dipole moment means how polar the molecule is, and since the hydrogen halides are diatomic and linear, this translates directly into how polar the bond is. The polarity of a bond depends on the electronegativity difference between the 2 atoms. The general trend in electronegativity is that it decreases down a column. This is because, while elements in the same column have the same effective nuclear charge (+7 in the case of halogens), they have larger electron clouds as you move down the periodic table. This means the valence (bonded) electrons are farther from the nucleus of Br than they are from the nucleus of F. Therefore, the attraction they feel for the nucleus is less for the larger atoms.

Complication Dipole moment decrease down the group so why HBr have boiling points ?? It is known that the dipole moment of HCl is greater than that of HI. If we evaluate this based on dipole moment alone then the boiling point of HCl might be thought to be higher than that of HI, but it is not. In the same family, HF is known to have a much higher boiling point because of hydrogen bonding but this effect is not present in the remaining molecules, HCl, HBr, and HI. It is known that molar mass (number of electrons/ number of electron shells) is a contributor to the property of boiling point. As molar mass increases the boiling point of similar molecules is observed to increase.

So what is the reason that HI has the highest boiling point of the three molecules? . It is the polarizability of the molecules that has a major effect on the boiling point of a series of molecules like HCl, HBr, and HI. Since HI is much more polarizable than the other molecules in the series Polarizability indicates how readily an electron cloud can be distorted. The electrons which are the most easily displaced in an atom or molecule are the valence electrons, these are the furthest from the nucleus. So valence electrons make the greatest contribution to the polarizability. The larger the atom, the more loosely it holds the electrons in its outermost shell, and the more they can be distorted. The more polarizable or the easier it is to create instantaneous dipoles. the stronger are the van der Waals interactions so grater the boiling point So the increase in London forces (from HCl to HI) outweighs the decrease in permanent dipole

More about polaribility
The polarizability of N2 (non-polar) is greater than H2(non-polar) , and that of CCl4(nonpolar) greater than CH4 (non-polar), and CO2(non-polar) greater than that of CO(polar). The strength of the intermolecular attractive forces is reflected in the boiling points of the substances. More polaribility in non-polar substances means more Van der Waals forces Although CO2 has more van der Waals forces CO has higher boiling point as it has stronger dipole-dipole forces as it is polar

Boiling temperature between HF and Water ?
 Water forms (up to) two hydrogen bonds (per molecule but HF only One)  So more energy is required to separate water molecule than HF  Water has a higher boiling temperature than HF

Rules for predicting whether a molecule is polar (has a permanent dipole) or is nonpolar :
If the central atom has one or more unshared, nonbonding, pairs of electrons the molecule is most likely polar. Examples include; NH3, H2O, SO2 If the central atom has no unshared, or nonbonding, pairs of electrons and nonidentical terminal atoms the molecule is polar. If the central atom has has no unshared, or nonbonding, pairs of electrons and the terminal atoms are identical the molecule is nonpolar Leniar molecules like CO arepolar. Molecules with single

Dipole-Dipole

Hybridization-The Blending of Orbitals.
Dipole- is created by equal but opposite charges that are separated by a short distance Dipole-Dipole Attractions-Attraction between oppositely charged regions of neighboring molecules. Hydrogen Bonding- Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen. Hydrogen bonding in Kevlar, a strong polymer used in bullet-proof vests London Dispersion Forces- The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. London forces increase with the size of the molecules.

Affects of Intermolecular Forces
Solubility Vapor Pressures Freezing Temperature Boiling Temperature Surface tension

Comparison of IMF’s Ion-Dipole > H-bonding > dipole-dipole > dispersion

Solubility LIKE DISSOVES LIKE
polar compounds tend to dissolve in polar solvents (like water), and that non-polar compounds tend to dissolve in non-polar solvents polar compounds tend to dissolve in polar solvents (like water because they can form polar bonds with the solvent (hydrogen bonds , dipole-dipole interactions) non-polar compounds does not dissolve in polar solvents because non polar compounds have weaker london forces which do not have enough energy to break polar bonds like hydrogen bonds, dipole-dipole bonds

Non polar, so cannot break
strong hydrogen bondsbetween water molecules Non polar (As symmetrical molecule), so cannot break strong hydrogen bondsbetween water molecules

2 Polar substances dissolves

Solubility Highly polar solids such as ionic salts such (example: Sodium cloride) dissove in water ( a polar solvent) but not dissove in hexane (non polar solvent) Polar organic substances dissove in water but does not dissove in hexane Non polar liquids such as candle wax does not dissove in water but dissove in hexane Non-polar liquids such as petrol and diesel mix completely Polar liquids such as ethanol and water dissoves

An alcohol has both a nonpolar alkyl group and a polar OH group
An alcohol has both a nonpolar alkyl group and a polar OH group. So is an alcohol molecule nonpolar or polar? Is it soluble in a nonpolar solvent, or is it soluble in water? The answer depends on the size of the alkyl group. As the alkyl group increases in size, it becomes a more significant fraction of the alcohol molecule and the compound becomes less and less soluble in water. In other words, the molecule becomes more and more like an alkane. Four carbons tend to be the dividing line at room temperature. Alcohols with fewer than four carbons are soluble in water, but alcohols with more than four carbons are insoluble in water. In other words, an OH group can drag about three or four carbons into solution in water.  Alcohols with branched alkyl groups are more soluble in water than alcohols with nonbranched alkyl groups with the same number of carbons, because branching minimizes the contact surface of the nonpolar portion of the molecule.

Solubility of halogenoalkanes Solubility in water
The halogenoalkanes are at best only very slightly soluble in water. In order for a halogenoalkane to dissolve in water you have to break attractions between the halogenoalkane molecules (van der Waals dispersion and dipole-dipole interactions) and break the hydrogen bonds between water molecules. Both of these cost energy. Energy is released when new attractions are set up between the halogenoalkane and the water molecules. These will only be dispersion forces and dipole-dipole interactions. These aren't as strong as the original hydrogen bonds in the water, and so not as much energy is released as was used to separate the water molecules. The energetics of the change are sufficiently "unprofitable" that very little dissolves.

Solubility of halogenoalkanes in organic solvents
Halogenoalkanes tend to dissolve in organic solvents because the new intermolecular attractions have much the same strength as the ones being broken in the separate halogenoalkane and solvent.

Solubility Ionic solids in water
In ionic solids, ions are held by strong electrostatic attractions between positive and negatively charged ions Energy to break own ionic lattice is known as Lattice enthalpy Enthalpy of hydration, Hhyd, of an ion is the enthalpy change when 1 mole of gaseous ions dissolve in sufficient water to give an infinitely dilute solution. Hydration enthalpies are always negative. If lattice energy > hydration enthalpy the substance is not soluable If lattice energy < hydration enthalpy the substance is dissoves (exorthermic process) If lattice energy = hydration enthalpy enthalpy the substance is dissoves

Testing concepts Which attractions are stronger: intermolecular or intramolecular? How many times stronger is a covalent bond compared to a dipole-dipole attraction? What evidence is there that nonpolar molecules attract each other? Which chemical in table 10.1 has the weakest intermolecular forces? Which has the strongest? How can you tell? Suggest some ways that the dipoles in London forces are different from the dipoles in dipole-dipole attractions. A) Which would have a lower boiling point: O2 or F2? Explain. B) Which would have a lower boiling point: NO or O2? Explain.

Which would you expect to have the higher melting point (or boiling point): C8H18 or C4H10? Explain.
What two factors causes hydrogen bonds to be so much stronger than typical dipole-dipole bonds? So far we have discussed 4 kinds of intermolecular forces: ionic, dipole-dipole, hydrogen bonding, and London forces. What kind(s) of intermolecular forces are present in the following substances: a) NH3, b) SF6, c) PCl3, d) LiCl, e) HBr, f) CO2 (hint: consider EN and molecular shape/polarity) Challenge: Ethanol (CH3CH2OH) and dimethyl ether (CH3OCH3) have the same formula (C2H6O). Ethanol boils at 78 C, whereas dimethyl ether boils at -24 C. Explain why the boiling Temperature of the ether is so much lower than the boiling point of ethanol. Challenge: try answering the question on the next slide.

Testing concepts Intramolecular are stronger.
A covalent bond is 100x stronger. The molecules gather together as liquids or solids at low temperatures. Based on boiling points, F2 (-188) has the weakest forces, H2S has the strongest (-61). London forces Are present in all compounds Can occur between atoms or molecules Are due to electron movement not to EN Are transient in nature (dipole-dipole are more permanent). London forces are weaker

Testing concepts A) F2 would be lower because it is smaller. Larger atoms/molecules can have their electron clouds more easily deformed and thus have stronger London attractions and higher melting/boiling points. B) O2 because it has only London forces. NO has a small EN, giving it small dipoles. C8H18 would have the higher melting/boiling point. This is a result of the many more sites available for London forces to form. 1) a large EN, 2) the small sizes of atoms.

Testing concepts a) NH3: Hydrogen bonding (H + N), London.
b) SF6: London only (it is symmetrical). c) PCl3: EN= Dipole-dipole, London. d) LiCl: EN= Ionic, (London). e) HBr: EN= Dipole-dipole, London. f) CO2: London only (it is symmetrical) Challenge: In ethanol, H and O are bonded (the large EN results in H-bonding). In dimethyl ether the O is bonded to C (a smaller EN results in a dipole-dipole attraction rather than hydrogen bonding).

Of the following substances
PH3 CH H2O CO2 SO2 What is the predominant intermolecular force in each substance? Which has the lowest heat of vaporization? Which is the best example of H-bonding? Which is often used as a supercritical fluid? Which should be the best solvent for NH4Cl?

Intermolecular forces- IMF

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