Intermolecular Forces Chemical and Physical Forces Involved in Chemical Bond Formation and Molecules, Mechanisms of Bond Formati

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Intermolecular Forces Chemical and physical forces involved in chemical bond formation and molecules, mechanisms of bond formation pi-bond, covalent, electrostatic, co-ordinate bond, hydrophobic interaction and their properties

Mitesh Shrestha Bonds  Forces that hold groups of atoms together and make them function as a unit. Example: H-O-H Bond Energy: Energy required to break a bond. Ionic Bond: Attractions between oppositely charged ions. Example: Na+ Cl-

Bonding Forces

 Electron – electron repulsive forces  Nucleus – nucleus repulsive forces  Electron – nucleus attractive forces

2 Chemical Bonds Chemical Bonds • a strong force of attraction holding atoms together in a molecule or crystal, resulting from the sharing or transfer of electrons. Types of chemical bonds

Ionic Compound: A compound resulting from a positive ion (usually a metal) combining with a negative ion (usually a non-metal). Example: M+ + X-  MX Covalent Bond: Electrons are shared by nuclei. Example: H-H Polar Covalent Bond: Unequal sharing of electrons by nuclei. Example: H-F Hydrogen fluoride is an example of a molecule that has bond polarity. Types of Chemical Bonding

1. Metal with nonmetal:

electron transfer and ionic bonding

6 Three models of chemical bonding Ionic

Electron transfer 7 IONic Bonding

• electrons are transferred between valence shells of atoms • ionic compounds are made of ions NOT MOLECULES

• ionic compounds are called Salts or Crystals

IONic bonding

• Always formed between metals and non- metals

- [METALS ]+ [NON-METALS ]

- Lost e Gained e- IONic Bonding

• Electronegativity difference > 2.0 – Look up e-neg of the atoms in the bond and subtract NaCl

CaCl2 • Compounds with polyatomic ions

NaNO3

Properties of Ionic Compounds

SALTS • hard solid @ 22oC Crystals • high mp temperatures • nonconductors of electricity in solid phase • good conductors in liquid phase or dissolved in water (aq) Types of Chemical Bonding

1. Metal with nonmetal:

electron transfer and ionic bonding

2. Nonmetal with nonmetal:

electron sharing and covalent bonding

13 Three models of chemical bonding

Ionic Covalent

Electron transfer Electron sharing 14 Covalent Bonding

• Pairs of e- are shared molecules between non-metal atoms • electronegativity difference < 2.0 • forms polyatomic ions Properties of Molecular Substances

Covalent bonding

• Low m.p. temp and b.p. temps • relatively soft solids as compared to ionic compounds • nonconductors of electricity in any phase Types of Covalent Bonds • NON-Polar bonds –Electrons shared evenly in the bond –E-neg difference is zero

Between identical atoms Diatomic molecules Types of Covalent Bonds

Polar bond –Electrons unevenly shared –E-neg difference greater than zero but less than 2.0

closer to 2.0 more polar more “ionic character” Polar Covalent Bond

• Unequal sharing of electrons between atoms in a molecule. • One atom attracts the electrons more than the other atom. • Results in a charge separation in the bond (partial positive and partial negative charge).

19 Polar Molecules

• Molecules with a positive and a negative end • Requires two things to be true  The molecule must contain polar bonds This can be determined from differences in electronegativity. Symmetry can not cancel out the effects of the polar bonds. Must determine geometry first. How to show a bond is polar • Isn’t a whole charge just a partial charge • d+ means a partially positive • d- means a partially negative

d+ d-

H Cl • The Cl pulls harder on the electrons • The electrons spend more time near the Cl Electronegativity

• A measure of how strongly the atoms attract electrons in a bond. • The bigger the electronegativity difference the more polar the bond. • 0.0 - 0.3 Covalent nonpolar • 0.3 - 1.67 Covalent polar • >1.67 Ionic

Electronegativity

• The ability of an atom in a molecule to attract shared electrons to itself. • For a molecule HX, the relative electronegativities of the H and X atoms are determined by comparing the measured H–X bond energy with the “expected” H–X bond energy.

23 Electronegativity

• On the periodic table, electronegativity generally increases across a period and decreases down a group. • The range of electronegativity values is from 4.0 for fluorine (the most electronegative) to 0.7 for cesium and francium (the least electronegative).

24 Electronegativity Values for Selected Elements

• The polarity of a bond depends on the difference between the electronegativity values of the atoms forming the bond.

1. Metal with nonmetal:

electron transfer and ionic bonding

2. Nonmetal with nonmetal:

electron sharing and covalent bonding

3. Metal with metal:

electron pooling and metallic bonding

27 Three models of chemical bonding

Ionic Covalent Metallic

Electron transfer Electron sharing Electron pooling 28 Metallic bonding

• Metallic bonding – Occurs between like atoms of a metal in the free state – Valence e- are mobile (move freely among all metal atoms) – Positive ions in a sea of electrons

• Metallic characteristics – High mp temps, ductile, malleable, shiny – Hard substances – Good conductors of heat and electricity as (s) and (l) Electrostatic (Lattice) Energy

Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions.

Q+ is the charge on the cation Q Q E = k + - r Q- is the charge on the anion r is the distance between the ions

cmpd lattice energy

MgF2 2957 Q= +2,-1

Lattice energy (E) increases as Q increases MgO 3938 Q= +2,-2 and/or as r decreases. LiF 1036 r F < r Cl LiCl 853 30 Coordinate Covalent Bond

• When one atom donates both electrons in a covalent bond. • Carbon monoxide • CO C O Coordinate Covalent Bond  When one atom donates both electrons in a covalent bond.  Carbon monoxide  CO C O Coordinate Covalent Bond  When one atom donates both electrons in a covalent bond.  Carbon monoxide  CO C O

Pi Bonds

• Pi (π) bonds are formed when the bonding electron pair is placed in a molecular orbital formed (frequently) by p orbitals on adjacent atoms overlapping. • The p orbitals on the bonded atoms are oriented perpendicular to the internuclear axis (which makes orbital overlap slightly less favourable).

Physical bonds

• Van der Waals force • London dispersion force (part of the van der Waals forces) • Dipole–dipole interactions • Debye force (induced dipole) • Hydrogen bonding

Van der Waals

• The residual attractive or repulsive forces between molecules or atomic groups that do not arise from a covalent bond, or electrostatic interaction of ions or of ionic groups with one another or with neutral molecules. • The resulting van der Waals forces can be attractive or repulsive. Van der Waals • They are weaker than normal covalent and ionic bonds. • Van der Waals forces are additive and cannot be saturated. • They have no directional characteristic. • They are all short-range forces and hence only interactions between the nearest particles need to be considered (instead of all the particles). Attraction is greater if the molecules are closer, due to Van der Waals forces. • Van der Waals forces are independent of temperature except dipole - dipole interactions.

Van der Waals

• Non-polar molecules can exist in liquid and solid phases because van der Waals forces keep the molecules attracted to each other

• Exist between CO2, CH4, CCl4, CF4, diatomics and monoatomics Van der Waals periodicity

• increase with molecular mass.

• increase with closer distance between molecules – Decreases when particles are farther away

London dispersion force • A type of force acting between atoms and molecules. • Exhibited by nonpolar molecules because of the correlated movements of the electrons in interacting molecules. Because the electrons in adjacent molecules "flee" as they repel each other, electron density in a molecule becomes redistributed in proximity to another molecule. • This is frequently described as the formation of instantaneous dipoles that attract each other. London forces are present between all chemical groups, and usually represent the main part of the total interaction force in condensed matter, even though they are generally weaker than ionic bonds and hydrogen bonds.

London dispersion force • The London dispersion force is the weakest intermolecular force. • The London dispersion force is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles. • This force is sometimes called an induced dipole- induced dipole attraction. London forces are the attractive forces that cause nonpolar substances to condense to liquids and to freeze into solids when the temperature is lowered sufficiently. London dispersion force • Because of the constant motion of the electrons, an atom or molecule can develop a temporary (instantaneous) dipole when its electrons are distributed unsymmetrically about the nucleus.

London dispersion force • A second atom or molecule, in turn, can be distorted by the appearance of the dipole in the first atom or molecule (because electrons repel one another) which leads to an electrostatic attraction between the two atoms or molecules.

• Dispersion forces are present between any two molecules (even polar molecules) when they are almost touching.

London dispersion force • Molecular Size – Dispersion forces are present between all molecules, whether they are polar or nonpolar. – Larger and heavier atoms and molecules exhibit stronger dispersion forces than smaller and lighter ones. – In a larger atom or molecule, the valence electrons are, on average, farther from the nuclei than in a smaller atom or molecule. They are less tightly held and can more easily form temporary dipoles. – The ease with which the electron distribution around an atom or molecule can be distorted is called the polarizability. • London dispersion forces tend to be: – stronger between molecules that are easily polarized. – weaker between molecules that are not easily polarized.

London dispersion force • Molecular Shape – The shapes of molecules also affect the magnitudes of dispersion forces between them.

– At room temperature, neopentane (C5H12) is a gas whereas n-pentane (C5H12) is a liquid. – London dispersion forces between n-pentane molecules are stronger than those between neopentane molecules even though both molecules are nonpolar and have the same molecular weight. – The somewhat cylindrical shape of n-pentane molecules allows them to come in contact with each other more effectively than the somewhat spherical neopentane molecules.

Intermolecular Forces

What holds molecules to each other Intermolecular Forces • The forces with which molecules attract each other.

• Intermolecular forces are weaker than ionic or covalent bonds.

• Intermolecular forces are responsible for the physical state of a compound (solid, liquid or gas).

Intermolecular Forces • The weakest are called van der Waal’s forces - there are two kinds – Dispersion forces – Dipole Interactions • depend on the number of electrons • more electrons stronger forces • Bigger molecules Dipole • A dipole is a molecule that has both positive and negative regions. • Molecular dipoles occur due to the unequal sharing of electrons between atoms in a molecule. • Those atoms that are more electronegative pull the bonded electrons closer to themselves. • The buildup of electron density around an atom or discreet region of a molecule can result in a molecular dipole in which one side of the molecule possesses a partially negative charge and the other side a partially positive charge. • Molecules with dipoles that are not canceled by their molecular geometry are said to be polar. Dipole Moment

• Property of a molecule whose charge distribution can be represented by a center of positive charge and a center of negative charge. • Use an arrow to represent a dipole moment. . Point to the negative charge center with the tail of the arrow indicating the positive center of charge. Dipole Moment in a Water Molecule • The polarity of water affects its properties. . Permits ionic compounds to dissolve in it.

. Causes water to remain liquid at higher temperature. Table 13.1 Dipole–dipole interactions (Keesom force) • Dipole-dipole forces are attractive forces between the positive end of one polar molecule and the negative end of another polar molecule. • Dipole-dipole forces have strengths that range from 5 kJ to 20 kJ per mole. • They are much weaker than ionic or covalent bonds and have a significant effect only when the molecules involved are close together (touching or almost touching).

Dipole–dipole interactions • Polar molecules have a partial negative end and a partial positive end. • The partially positive end of a polar molecule is attracted to the partially negative end of another. • In a ICl molecule the more electronegative chlorine atom bears the partial negative charge; the less electronegative iodine atom bears the partial positive charge. • The partially positive iodine end of one ICl molecule is attracted to the partially negative chlorine end of another ICl molecule. • A dashed line is used to represent an intermolecular attraction between molecules because these forces are NOT as strong as chemical bonds.

Dipole–dipole interactions

• Depend on the number of electrons • More electrons stronger forces • Bigger molecules more electrons • Fluorine is a gas • Bromine is a liquid • Iodine is a solid Dipole–dipole interactions

• Occur when polar molecules are attracted to each other.

• Slightly stronger than dispersion forces.

• Opposites attract but not completely hooked like in ionic solids. Dipole–dipole interactions • Occur when polar molecules are attracted to each other.

• Slightly stronger than dispersion forces.

• Opposites attract but not completely hooked like in ionic solids. d+ d- d+ d- H F H F Dipole–dipole interactions

d+ d- d+ d- Biological Importance of Dipole Interactions • The potential energy from dipole interactions is important for living organisms. The biggest impact dipole interactions have on living organisms is seen with protein folding. Every process of protein formation, from the binding of individual amino acids to secondary structures to tertiary structures and even the formation of quaternary structures is dependent on dipole-dipole interactions. • A prime example of quaternary dipole interaction that is vital to human health is the formation of erythrocytes. Erythrocytes, commonly known as red blood cells, are comprised of four protein subunits and a heme molecule. For an erythrocyte to form properly, multiple steps must occur, all of which involve dipole interactions. The four protein subunits—two alpha chains, two beta chains—and the heme group, interact with each other through a series of dipole-dipole interactions which allow the erythrocyte to take its final shape. Any mutation that destroys these dipole-dipole interactions prevents the erythrocyte from forming properly, and impairs their ability to carry oxygen to the tissues of the body. So we can see that without the dipole-dipole interactions, proteins would not be able to fold properly and all life as we know it would cease to exist.

Ion – dipole interaction • An ion-dipole force consists of an ion and a polar molecule interacting. They align so that the positive and negative groups are next to one another, allowing for maximum attraction.

Ion - Dipole Forces

• Ion - dipole forces exist between ions and polar molecules.

• The magnitude of these forces increases as: –the distance between the ion and the polar molecule decreases –the magnitude of the charge on the ion increases –the magnitude of the dipole of the polar molecule increases. Ion - Dipole Forces • Hydration energies for cations and anions is an excellent example of this concept.

• When these hydration bond form, energy is released, exothermic.

• This energy is then used to break the ion - ion forces in the ionic solid.

• When the hydration energy is large enough, the ionic solid is soluble in water. Attraction Between Ions and Permanent Dipoles •• - d water

O dipole ••

H H + d Water is highly polar and can interact with positive ions to give hydrated ions in water. Induced-Dipole Forces

• An ion-induced dipole attraction is a weak attraction that results when the approach of an ion induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.

Ion- Induced Dipole Dipole-Induced Dipole Forces (Debye force)

• A dipole-induced dipole attraction is a weak attraction that results when a polar molecule induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species. Dipole-Induced Dipole Forces

• What would happen if we mixed HCl with argon, which has no dipole moment? The electrons on an argon atom are distributed homogeneously around the nucleus of the atom. But these electrons are in constant motion. When an argon atom comes close to a polar HCl molecule, the electrons can shift to one side of the nucleus to produce a very small dipole moment that lasts for only an instant.

• By distorting the distribution of electrons around the argon atom, the polar HCl molecule induces a small dipole moment on this atom, which creates a weak dipole-induced dipole force of attraction between the HCl molecule and the Ar atom. This force is very weak, with a bond energy of about 1 kJ/mol

Hydrogen Bonds

• It is a force of attraction between a hydrogen atom in one molecule and a small atom of high electronegativity in another molecule. That is, it is an intermolecular force, not an intramolecular force as in the common use of the word bond.

• When hydrogen atoms are joined in a polar covalent bond with a small atom of high electronegativity such as O, F or N, the partial positive charge on the hydrogen is highly concentrated because of its small size. If the hydrogen is close to another oxygen, fluorine or nitrogen in another molecule, then there is a force of attraction termed a dipole-dipole interaction. This attraction or "hydrogen bond" can have about 5% to 10% of the strength of a covalent bond.

• Hydrogen bonding has a very important effect on the properties of water and ice. Hydrogen bonding is also very important in proteins and nucleic acids and therefore in life processes. The "unzipping" of DNA is a breaking of hydrogen bonds which help hold the two strands of the double helix together. Hydrogen bonding

• Are the attractive force caused by hydrogen bonded to F, O, or N. • F, O, and N are very electronegative so it is a very strong dipole. • The hydrogen partially share with the lone pair in the molecule next to it. • The strongest of the intermolecular forces. H is shared between 2 atoms of OXYGEN or 2 atoms of NITROGEN or 2 atoms of FLUORINE Of 2 different molecules Why does H “bonding” occur?

• Nitrogen, Oxygen and Fluorine – small atoms with strong nuclear charges • powerful atoms – very high electronegativities

Hydrogen Bonding

d+ d- H O H d+ Hydrogen bonding

H O H Hydrogen Bonding

Hydrogen bonding and base pairing in DNA

Adenine

Thymine

79 Figure 13.14 WHAT IS MEANT HYDROPHOBICITY ? Hydrophobicity comes also from the greek word Hydro(water) and Phobicity (fear) it refers to the physical property of a material that repels a mass of water.

• Some of the common natural Hydrophobic materials are waxes, oil and fats.

The evaluation of hydrophobicity is made through water contact angle measurements.

A water droplet would be spherical so the water contact angle will be significantly high.

HYDROPHOBIC EFFECT Continued….

The hydrophobic effect represents the tendency of water to exclude nonpolar molecules. This occurs because interactions between the hydrophobic molecules enable the displaced water molecules to make hydrogen bonds more freely with each other and increase the number of hydrogen bonds they are involved with, thereby decreasing the overall free energy. The word hydrophobic literally means "water-fearing," and it describes the segregation and apparent repulsion between water and nonpolar substances.

• At the molecular level, the hydrophobic effect is important in driving protein folding formation of lipid bilayers and micelles, insertion of membrane proteins into the nonpolar lipid environment and protein- small molecule interactions. Substances for which this effect is observed are known as hydrophobes. SUPER HYDROPHOBIC COATING

The process of coating the surface of a material with hydrophobic property material in order to avoid sticking of liquids on that surface.

This is absolutely unique way of coating unlike conventional which shrink continuously during drying to produce low porosity films.

Super hydrophobic technology makes water bounce, it stops it, rolls it off the surface. Young's equation is used to describe the interactions between the forces of cohesion and adhesion and measure what is referred to as surface energy.

www.ramehart.com/contactangle.htm Continued.. • A drop with a contact angle over 90° is hydrophobic.

• This condition is exemplified by poor wetting, poor adhesiveness and the solid surface free energy is low.

Hydrophobic interactions

• Hydrophobic interactions describe the relations between water and hydrophobes (low water-soluble molecules). • Hydrophobes are nonpolar molecules and usually have a long chain of carbons that do not interact with water molecules. • The mixing of fat and water is a good example of this particular interaction. • The common misconception is that water and fat doesn’t mix because the Van der Waals • forces that are acting upon both water and fat molecules are too weak. Causes of Hydrophobic Interactions

• American chemist Walter Kauzmann discovered that nonpolar substances like fat molecules tend to clump up together rather that distributing itself in a water medium, because this allow the fat molecules to have minimal contact with water.

Hydrophobic interactions - ChemWiki Formation of Hydrophobic Interactions

• The mixing hydrophobes and water molecules is not spontaneous; however, hydrophobic interactions between hydrophobes are spontaneous.

• When hydropobes come together and interact with each other, enthalpy increases ( is positive) because some of hydrogen bonds that form the clathrate cage will be broken.

• Tearing down a portion of the clathrate cage will cause the entropy to increase ( is positive), since forming it decreases the entropy. According to the formula: ΔG= ΔH-TΔS

ΔH=Small positive value ΔS=Large positive value

Result :ΔG= Negative

A negative ΔG indicates that hydrophobic interactions are spontaneous

Strength of Hydrophobic Interactions

• Temperature: As temperature increases, the strength of hydrophobic interactions increases also. However, at an extreme temperature, hydrophobic interactions will denature.

• Number of carbons on the hydrophobes: Molecules with the greatest number of carbons will have the strongest hydrophobic interactions.

• The shape of the hydrophobes: Aliphatic organic molecules have stronger interactions than aromatic compounds. Branches on a carbon chain will reduce the hydrophobic effect of that molecule and linear carbon chain can produce the largest hydrophobic interaction. This is so because carbon branches produce steric hindrance, so it is harder for two hydrophobes to have very close interactions with each other to minimize their contact to water.

Biological Importance of Hydrophobic Interactions • Hydrophobic Interactions are important for the folding of proteins. This is important in keeping a protein alive and biologically active, because it allows the protein to decrease in surface area and reduce the undesirable interactions with water. • Besides from proteins, there are many other biological substances that rely on hydrophobic interactions for its survival and functions, like the phospholipid bilayer membranes in every cell of your body. APPLICATIONS

• A primary purpose of hydrophobic coatings such as polytetrafluoroethylene(PTFE) or polyxylylene is to act as a barrier against water commonly seen in automobiles

• Used in fabrication on metallic nano rod to prevent icing.

• Its is widely used in aerospace industry for providing anti-icing coating on the surface of the aeroplane .

• Hydrophobic self cleaning glasses are installed in traffic sensor control unit.

• We induce hydrophobic recovery after plasma treatment, a physical contact treatment (PCT) .