van der Waals force

From Wikipedia, the free encyclopedia

https://en.wikipedia.org/wiki/Van_der_Waals_force

Rainwater flux from a canopy. Among the forces that govern drop formation: Van der Waals force, surface tension, cohesion, Plateau–Rayleigh instability.

Microfiber cloth makes use of van der Waals force to remove dirt without scratches.

In molecular physics and chemistry, the van der Waals force is a distance-dependent interaction between atoms or molecules. Unlike ionic or covalent bonds, these attractions do not result from a chemical electronic bond; they are comparatively weak and therefore more susceptible to disturbance. The van der Waals force quickly vanishes at longer distances between interacting molecules.

Named after Dutch physicist Johannes Diderik van der Waals, the van der Waals force plays a fundamental role in fields as diverse as supramolecular chemistry, structural biology, polymer science, nanotechnology, surface science, and condensed matter physics. It also underlies many properties of organic compounds and molecular solids, including their solubility in polar and non-polar media.

If no other force is present, the distance between atoms at which the force becomes repulsive rather than attractive as the atoms approach one another is called the van der Waals contact distance; this phenomenon results from the mutual repulsion between the atoms' electron clouds.

The van der Waals forces are usually described as a combination of the London dispersion forces between "instantaneously induced dipoles", Debye forces between permanent dipoles and induced dipoles, and the Keesom force between permanent molecular dipoles whose rotational orientations are dynamically averaged over time.

Definition

Van der Waals forces include attraction and repulsions between atoms, molecules, as well as other intermolecular forces. They differ from covalent and ionic bonding in that they are caused by correlations in the fluctuating polarizations of nearby particles (a consequence of quantum dynamics).

The force results from a transient shift in electron density. Specifically, the electron density may temporarily shift to be greater on one side of the nucleus. This shift generates a transient charge which a nearby atom can be attracted to or repelled by. The force is repulsive at very short distances, reaches zero at an equilibrium distance characteristic for each atom, or molecule, and becomes attractive for distances larger than the equilibrium distance. For individual atoms, the equilibrium distance is between 0.3 nm and 0.5 nm, depending on the atomic-specific diameter. When the interatomic distance is greater than 1.0 nm the force is not strong enough to be easily observed as it decreases as a function of distance r approximately with the 7th power (~r⁻⁷).

Van der Waals forces are often among the weakest chemical forces. For example, the pairwise attractive van der Waals interaction energy between H (Hydrogen) atoms in different H₂ molecules equals 0.06 kJ/mol (0.6 meV) and the pairwise attractive interaction energy between O (Oxygen) atoms in different O₂ molecules equals 0.44 kJ/mol (4.6 meV). The corresponding vaporization energies of H₂ and O₂ molecular liquids, which result as a sum of all van der Waals interactions per molecule in the molecular liquids, amount to 0.90 kJ/mol (9.3 meV) and 6.82 kJ/mol (70.7 meV), respectively, and thus approximately ~15 times the value of the individual pairwise interatomic interactions (excluding covalent bonds).

The strength of van-der-Waals bonds increases with higher polarizability of the participating atoms. For example, the pairwise van der Waals interaction energy for more polarizable atoms such as S (Sulfur) atoms in H₂S and sulfides exceeds 1 kJ/mol (10 meV), and the pairwise interaction energy between even larger, more polarizable Xe (Xenon) atoms is 2.35 kJ/mol (24.3 meV). These van der Waals interactions are up to 40 times stronger than in H₂, which has only one valence electron, and they are still not strong enough to achieve an aggregate state other than gas for Xe under standard conditions. The interactions between atoms in metals can also be effectively described as van-der-Waals interactions and account for the observed solid aggregate state with bonding strengths comparable to covalent and ionic interactions. The strength of pairwise van-der-Waals type interactions is on the order of 12 kJ/mol (120 meV) for low-melting Pb (Lead) and on the order of 32 kJ/mol (330 meV) for high-melting Pt (Platinum), which is about one order of magnitude stronger than in Xe due to the presence of a highly polarizable free electron gas. Accordingly, van der Waals forces can range from weak to strong interactions, and support integral structural loads when multitudes of such interactions are present.

More broadly, intermolecular forces have several possible contributions:

1. A repulsive component resulting from the Pauli exclusion principle that prevents close contact of atoms, or the collapse of molecules.
2. Attractive or repulsive electrostatic interactions between permanent charges (in the case of molecular ions), dipoles (in the case of molecules without inversion centre), quadrupoles (all molecules with symmetry lower than cubic), and in general between permanent multipoles. These interactions also include hydrogen bonds, cation-pi, and pi-stacking interactions. Orientation-averaged contributions from electrostatic interactions are sometimes called the Keesom interaction or Keesom force after Willem Hendrik Keesom.
3. Induction (also known as polarization), which is the attractive interaction between a permanent multipole on one molecule with an induced multipole on another. This interaction is sometimes called Debye force after Peter J.W. Debye.
4. Dispersion (usually named London dispersion interactions after Fritz London), which is the attractive interaction between any pair of molecules, including non-polar atoms, arising from the interactions of instantaneous multipoles.

Hereby, different texts may refer to a different spectrum of interactions using the term "van der Waals force". Typically, contributions (1) and (4) are considered as van-der-Waals forces, excluding effects from permanent multipoles as described in (2) and from permanent polarization in (3). However, some texts describe the van der Waals force as the totality of forces, including repulsion; others mean all the attractive forces (and then sometimes distinguish van der Waals–Keesom, van der Waals–Debye, and van der Waals–London).

All intermolecular/van der Waals forces are anisotropic (except those between two noble gas atoms), which means that they depend on the relative orientation of the molecules. The induction and dispersion interactions are always attractive, irrespective of orientation, but the electrostatic interaction changes sign upon rotation of the molecules. That is, the electrostatic force can be attractive or repulsive, depending on the mutual orientation of the molecules. When molecules are in thermal motion, as they are in the gas and liquid phase, the electrostatic force is averaged out to a large extent because the molecules thermally rotate and thus probe both repulsive and attractive parts of the electrostatic force. Random thermal motion can disrupt or overcome the electrostatic component of the van der Waals force but the averaging effect is much less pronounced for the attractive induction and dispersion forces.

The Lennard-Jones potential is often used as an approximate model for the isotropic part of a total (repulsion plus attraction) van der Waals force as a function of distance.

Van der Waals forces are responsible for certain cases of pressure broadening (van der Waals broadening) of spectral lines and the formation of van der Waals molecules. The London–van der Waals forces are related to the Casimir effect for dielectric media, the former being the microscopic description of the latter bulk property. The first detailed calculations of this were done in 1955 by E. M. Lifshitz. A more general theory of van der Waals forces has also been developed.

The main characteristics of van der Waals forces are:

• They are weaker than normal covalent and ionic bonds.
• The Van der Waals forces are additive in nature, consisting of several individual interactions, 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). Van der Waals attraction is greater if the molecules are closer.
• Van der Waals forces are independent of temperature except for dipole-dipole interactions.

In low molecular weight alcohols, the hydrogen-bonding properties of their polar hydroxyl group dominate other weaker van der Waals interactions. In higher molecular weight alcohols, the properties of the nonpolar hydrocarbon chain(s) dominate and determine their solubility.

Van der Waals forces are also responsible for the weak hydrogen bond interactions between unpolarized dipoles particularly in acid-base aqueous solution and between biological molecules.

London dispersion force

Main article: London dispersion force

London dispersion forces, named after the German-American physicist Fritz London, are weak intermolecular forces that arise from the interactive forces between instantaneous multipoles in molecules without permanent multipole moments. In and between organic molecules the multitude of contacts can lead to larger contribution of dispersive attraction, particularly in the presence of heteroatoms. London dispersion forces are also known as 'dispersion forces', 'London forces', or 'instantaneous dipole–induced dipole forces'. The strength of London dispersion forces is proportional to the polarizability of the molecule, which in turn depends on the total number of electrons and the area over which they are spread. Hydrocarbons display small dispersive contributions, the presence of heteroatoms lead to increased LD forces as function of their polarizability, e.g. in the sequence RI>RBr>RCl>RF. In absence of solvents weakly polarizable hydrocarbons form crystals due to dispersive forces; their sublimation heat is a measure of the dispersive interaction.

Van der Waals forces between macroscopic objects

For macroscopic bodies with known volumes and numbers of atoms or molecules per unit volume, the total van der Waals force is often computed based on the "microscopic theory" as the sum over all interacting pairs. It is necessary to integrate over the total volume of the object, which makes the calculation dependent on the objects' shapes. For example, the van der Waals interaction energy between spherical bodies of radii R₁ and R₂ and with smooth surfaces was approximated in 1937 by Hamaker (using London's famous 1937 equation for the dispersion interaction energy between atoms/molecules as the starting point) by:

${\displaystyle \begin{array}{rl}& U(z;R_1,R_2)=-\frac{A}{6}\left(\frac{2R_1{R}_2}{z^2-(R_1+R_2{)}^2}+\frac{2R_1{R}_2}{z^2-(R_1-R_2{)}^2}+\ln \left[\frac{z^2-(R_1+R_2{)}^2}{z^2-(R_1-R_2{)}^2}\right]\right)\end{array}}$

(1)

where A is the Hamaker coefficient, which is a constant (~10⁻¹⁹ − 10⁻²⁰ J) that depends on the material properties (it can be positive or negative in sign depending on the intervening medium), and z is the center-to-center distance; i.e., the sum of R₁, R₂, and r (the distance between the surfaces): ${\displaystyle z=R_1+R_2+r}$.

The van der Waals force between two spheres of constant radii (R₁ and R₂ are treated as parameters) is then a function of separation since the force on an object is the negative of the derivative of the potential energy function,${\displaystyle F_{\mathrm{V}\mathrm{d}\mathrm{W}}(z)=-\frac{d}{dz}U(z)}$. This yields:

${\displaystyle F_{\mathrm{V}\mathrm{d}\mathrm{W}}(z)=-\frac{A}{6}\frac{64R_1^3{R}_2^{3}z}{[z^2-(R_1+R_2{)}^2{]}^2[z^2-(R_1-R_2{)}^2{]}^2}}$

(2)

In the limit of close-approach, the spheres are sufficiently large compared to the distance between them; i.e., ${\displaystyle r\ll R_1}$ or ${\displaystyle R_2}$, so that equation (1) for the potential energy function simplifies to:

${\displaystyle U(r;R_1,R_2)=-\frac{A{R}_1{R}_2}{(R_1+R_2)6r}}$

(3)

with the force:

${\displaystyle F_{\mathrm{V}\mathrm{d}\mathrm{W}}(r)=-\frac{A{R}_1{R}_2}{(R_1+R_2)6r^2}}$

(4)

The van der Waals forces between objects with other geometries using the Hamaker model have been published in the literature.

From the expression above, it is seen that the van der Waals force decreases with decreasing size of bodies (R). Nevertheless, the strength of inertial forces, such as gravity and drag/lift, decrease to a greater extent. Consequently, the van der Waals forces become dominant for collections of very small particles such as very fine-grained dry powders (where there are no capillary forces present) even though the force of attraction is smaller in magnitude than it is for larger particles of the same substance. Such powders are said to be cohesive, meaning they are not as easily fluidized or pneumatically conveyed as their more coarse-grained counterparts. Generally, free-flow occurs with particles greater than about 250 μm.

The van der Waals force of adhesion is also dependent on the surface topography. If there are surface asperities, or protuberances, that result in a greater total area of contact between two particles or between a particle and a wall, this increases the van der Waals force of attraction as well as the tendency for mechanical interlocking.

The microscopic theory assumes pairwise additivity. It neglects many-body interactions and retardation. A more rigorous approach accounting for these effects, called the "macroscopic theory" was developed by Lifshitz in 1956. Langbein derived a much more cumbersome "exact" expression in 1970 for spherical bodies within the framework of the Lifshitz theory while a simpler macroscopic model approximation had been made by Derjaguin as early as 1934. Expressions for the van der Waals forces for many different geometries using the Lifshitz theory have likewise been published.

Use by geckos and arthropods

Gecko climbing a glass surface

Further information: Arthropod adhesion

The ability of geckos – which can hang on a glass surface using only one toe – to climb on sheer surfaces has been for many years mainly attributed to the van der Waals forces between these surfaces and the spatulae, or microscopic projections, which cover the hair-like setae found on their footpads.

There were efforts in 2008 to create a dry glue that exploits the effect, and success was achieved in 2011 to create an adhesive tape on similar grounds (i.e. based on van der Waals forces). In 2011, a paper was published relating the effect to both velcro-like hairs and the presence of lipids in gecko footprints.

A later study suggested that capillary adhesion might play a role, but that hypothesis has been rejected by more recent studies.

A 2014 study has shown that gecko adhesion to smooth Teflon and polydimethylsiloxane surfaces is mainly determined by electrostatic interaction (caused by contact electrification), not van der Waals or capillary forces.

Among the arthropods, some spiders have similar setae on their scopulae or scopula pads, enabling them to climb or hang upside-down from extremely smooth surfaces such as glass or porcelain.

Non-covalent interaction

From Wikipedia, the free encyclopedia

https://en.wikipedia.org/wiki/Non-covalent_interaction

 

See also: Intermolecular force

In chemistry, a non-covalent interaction differs from a covalent bond in that it does not involve the sharing of electrons, but rather involves more dispersed variations of electromagnetic interactions between molecules or within a molecule. The chemical energy released in the formation of non-covalent interactions is typically on the order of 1–5 kcal/mol (1000–5000 calories per 6.02×10²³ molecules). Non-covalent interactions can be classified into different categories, such as electrostatic, π-effects, van der Waals forces, and hydrophobic effects.

Non-covalent interactions are critical in maintaining the three-dimensional structure of large molecules, such as proteins and nucleic acids. They are also involved in many biological processes in which large molecules bind specifically but transiently to one another (see the properties section of the DNA page). These interactions also heavily influence drug design, crystallinity and design of materials, particularly for self-assembly, and, in general, the synthesis of many organic molecules.

The non-covalent interactions may occur between different parts of the same molecule (e.g. during protein folding) or between different molecules and therefore are discussed also as intermolecular forces.

Electrostatic interactions

For a more detailed introduction to electrostatics, see Coulomb's law.

Ionic

Main article: Ionic bond

Scheme 1. Process of NaF formation -- example of an electrostatic interaction

Ionic interactions involve the attraction of ions or molecules with full permanent charges of opposite signs. For example, sodium fluoride involves the attraction of the positive charge on sodium (Na⁺) with the negative charge on fluoride (F⁻). However, this particular interaction is easily broken upon addition to water, or other highly polar solvents. In water ion pairing is mostly entropy driven; a single salt bridge usually amounts to an attraction value of about ΔG =5 kJ/mol at an intermediate ion strength I, at I close to zero the value increases to about 8 kJ/mol. The ΔG values are usually additive and largely independent of the nature of the participating ions, except for transition metal ions etc.

These interactions can also be seen in molecules with a localized charge on a particular atom. For example, the full negative charge associated with ethoxide, the conjugate base of ethanol, is most commonly accompanied by the positive charge of an alkali metal salt such as the sodium cation (Na⁺).

Hydrogen bonding

Main article: Hydrogen bond

Hydrogen-bonding-in-water

A hydrogen bond (H-bond), is a specific type of interaction that involves dipole–dipole attraction between a partially positive hydrogen atom and a highly electronegative, partially negative oxygen, nitrogen, sulfur, or fluorine atom (not covalently bound to said hydrogen atom). It is not a covalent bond, but instead is classified as a strong non-covalent interaction. It is responsible for why water is a liquid at room temperature and not a gas (given water's low molecular weight). Most commonly, the strength of hydrogen bonds lies between 0–4 kcal/mol, but can sometimes be as strong as 40 kcal/mol In solvents such as chloroform or carbon tetrachloride one observes e.g. for the interaction between amides additive values of about 5 kJ/mol. According to Linus Pauling the strength of a hydrogen bond is essentially determined by the electrostatic charges. Measurements of thousands of complexes in chloroform or carbon tetrachloride have led to additive free energy increments for all kind of donor-acceptor combinations.

Halogen bonding

Main article: Halogen bonding

Figure 1. Anionic Lewis base forming a halogen bond with electron-withdrawn bromine (Lewis acid)

Halogen bonding is a type of non-covalent interaction which does not involve the formation nor breaking of actual bonds, but rather is similar to the dipole–dipole interaction known as hydrogen bonding. In halogen bonding, a halogen atom acts as an electrophile, or electron-seeking species, and forms a weak electrostatic interaction with a nucleophile, or electron-rich species. The nucleophilic agent in these interactions tends to be highly electronegative (such as oxygen, nitrogen, or sulfur), or may be anionic, bearing a negative formal charge. As compared to hydrogen bonding, the halogen atom takes the place of the partially positively charged hydrogen as the electrophile.

Halogen bonding should not be confused with halogen–aromatic interactions, as the two are related but differ by definition. Halogen–aromatic interactions involve an electron-rich aromatic π-cloud as a nucleophile; halogen bonding is restricted to monatomic nucleophiles.

Van der Waals forces

Main article: Van der Waals forces

Van der Waals forces are a subset of electrostatic interactions involving permanent or induced dipoles (or multipoles). These include the following:

• permanent dipole–dipole interactions, alternatively called the Keesom force
• dipole-induced dipole interactions, or the Debye force
• induced dipole-induced dipole interactions, commonly referred to as London dispersion forces

Hydrogen bonding and halogen bonding are typically not classified as Van der Waals forces.

Dipole–dipole

Main article: Dipole-dipole interactions

Figure 2. Dipole–dipole interactions between two acetone molecules, with the partially negative oxygen atom interacting with the partially positive carbon atom in the carbonyl.

Dipole-dipole interactions are electrostatic interactions between permanent dipoles in molecules. These interactions tend to align the molecules to increase attraction (reducing potential energy). Normally, dipoles are associated with electronegative atoms, including oxygen, nitrogen, sulfur, and fluorine.

For example, acetone, the active ingredient in some nail polish removers, has a net dipole associated with the carbonyl (see figure 2). Since oxygen is more electronegative than the carbon that is covalently bonded to it, the electrons associated with that bond will be closer to the oxygen than the carbon, creating a partial negative charge (δ⁻) on the oxygen, and a partial positive charge (δ⁺) on the carbon. They are not full charges because the electrons are still shared through a covalent bond between the oxygen and carbon. If the electrons were no longer being shared, then the oxygen-carbon bond would be an electrostatic interaction.

${\displaystyle \stackrel{\delta +}{\text{H}}-\stackrel{\delta -}{\text{Cl}}\cdots \stackrel{\delta +}{\text{H}}-\stackrel{\delta -}{\text{Cl}}}$

Often molecules contain dipolar groups, but have no overall dipole moment. This occurs if there is symmetry within the molecule that causes the dipoles to cancel each other out. This occurs in molecules such as tetrachloromethane. Note that the dipole-dipole interaction between two individual atoms is usually zero, since atoms rarely carry a permanent dipole. See atomic dipoles.

Dipole-induced dipole

Main article: Debye force

A dipole-induced dipole interaction (Debye force) is due to the approach of a molecule with a permanent dipole to another non-polar molecule with no permanent dipole. This approach causes the electrons of the non-polar molecule to be polarized toward or away from the dipole (or "induce" a dipole) of the approaching molecule. Specifically, the dipole can cause electrostatic attraction or repulsion of the electrons from the non-polar molecule, depending on orientation of the incoming dipole. Atoms with larger atomic radii are considered more "polarizable" and therefore experience greater attractions as a result of the Debye force.

London dispersion forces

Main article: London dispersion forces

London dispersion forces are the weakest type of non-covalent interaction. In organic molecules, however, the multitude of contacts can lead to larger contributions, particularly in the presence of heteroatoms. They are also known as "induced dipole-induced dipole interactions" and present between all molecules, even those which inherently do not have permanent dipoles. Dispersive interactions increase with the polarizability of interacting groups, but are weakened by solvents of increased polarizability. They are caused by the temporary repulsion of electrons away from the electrons of a neighboring molecule, leading to a partially positive dipole on one molecule and a partially negative dipole on another molecule. Hexane is a good example of a molecule with no polarity or highly electronegative atoms, yet is a liquid at room temperature due mainly to London dispersion forces. In this example, when one hexane molecule approaches another, a temporary, weak partially negative dipole on the incoming hexane can polarize the electron cloud of another, causing a partially positive dipole on that hexane molecule. In absence of solvents hydrocarbons such as hexane form crystals due to dispersive forces ; the sublimation heat of crystals is a measure of the dispersive interaction. While these interactions are short-lived and very weak, they can be responsible for why certain non-polar molecules are liquids at room temperature.

π-effects

π-effects can be broken down into numerous categories, including π-π interactions, cation-π & anion-π interactions, and polar-π interactions. In general, π-effects are associated with the interactions of molecules with the π-systems of conjugated molecules such as benzene.

π–π interaction

Figure 3. Various ways that benzene can interact intermolecularly. Note, however, that the sandwich configuration is not a favorable interaction compared to displaced or edge-to-face

π–π interactions are associated with the interaction between the π-orbitals of a molecular system. The high polarizability of aromatic rings lead to dispersive interactions as major contribution to so-called stacking effects. These play a major role for interactions of nucleobases e.g. in DNA. For a simple example, a benzene ring, with its fully conjugated π cloud, will interact in two major ways (and one minor way) with a neighboring benzene ring through a π–π interaction (see figure 3). The two major ways that benzene stacks are edge-to-face, with an enthalpy of ~2 kcal/mol, and displaced (or slip stacked), with an enthalpy of ~2.3 kcal/mol. The sandwich configuration is not nearly as stable of an interaction as the previously two mentioned due to high electrostatic repulsion of the electrons in the π orbitals.

Cation–π and anion–π interaction

Figure 4

Cation–pi interactions involve the positive charge of a cation interacting with the electrons in a π-system of a molecule. This interaction is surprisingly strong (as strong or stronger than H-bonding in some contexts), and has many potential applications in chemical sensors. For example, the sodium ion can easily sit atop the π cloud of a benzene molecule, with C₆ symmetry (See figure 4).

Anion–π interactions are very similar to cation–π interactions, but reversed. In this case, an anion sits atop an electron-poor π-system, usually established by the placement of electron-withdrawing substituents on the conjugated molecule

Figure 5.

Polar–π

Polar–π interactions involve molecules with permanent dipoles (such as water) interacting with the quadrupole moment of a π-system (such as that in benzene (see figure 5). While not as strong as a cation-π interaction, these interactions can be quite strong (~1-2 kcal/mol), and are commonly involved in protein folding and crystallinity of solids containing both hydrogen bonding and π-systems. In fact, any molecule with a hydrogen bond donor (hydrogen bound to a highly electronegative atom) will have favorable electrostatic interactions with the electron-rich π-system of a conjugated molecule.

Hydrophobic effect

Main article: Hydrophobic effect

The hydrophobic effect is the desire for non-polar molecules to aggregate in aqueous solutions in order to separate from water. This phenomenon leads to minimum exposed surface area of non-polar molecules to the polar water molecules (typically spherical droplets), and is commonly used in biochemistry to study protein folding and other various biological phenomenon. The effect is also commonly seen when mixing various oils (including cooking oil) and water. Over time, oil sitting on top of water will begin to aggregate into large flattened spheres from smaller droplets, eventually leading to a film of all oil sitting atop a pool of water. However the hydrophobic effect is not considered a non-covalent interaction as it is a function of entropy and not a specific interaction between two molecules, usually characterized by entropy.enthalpy compensation. An essentially enthalpic hydrophobic effect materializes if a limited number of water molecules are restricted within a cavity; displacement of such water molecules by a ligand frees the water molecules which then in the bulk water enjoy a maximum of hydrogen bonds close to four.

Examples

Drug design

Most pharmaceutical drugs are small molecules which elicit a physiological response by "binding" to enzymes or receptors, causing an increase or decrease in the enzyme's ability to function. The binding of a small molecule to a protein is governed by a combination of steric, or spatial considerations, in addition to various non-covalent interactions, although some drugs do covalently modify an active site (see irreversible inhibitors). Using the "lock and key model" of enzyme binding, a drug (key) must be of roughly the proper dimensions to fit the enzyme's binding site (lock). Using the appropriately sized molecular scaffold, drugs must also interact with the enzyme non-covalently in order to maximize binding affinity binding constant and reduce the ability of the drug to dissociate from the binding site. This is achieved by forming various non-covalent interactions between the small molecule and amino acids in the binding site, including: hydrogen bonding, electrostatic interactions, pi stacking, van der Waals interactions, and dipole–dipole interactions.

Non-covalent metallo drugs have been developed. For example, dinuclear triple-helical compounds in which three ligand strands wrap around two metals, resulting in a roughly cylindrical tetracation have been prepared. These compounds bind to the less-common nucleic acid structures, such as duplex DNA, Y-shaped fork structures and 4-way junctions.

Protein folding and structure

Main article: Protein folding

The folding of proteins from a primary (linear) sequence of amino acids to a three-dimensional structure is directed by all types of non-covalent interactions, including the hydrophobic forces and formation of intramolecular hydrogen bonds. Three-dimensional structures of proteins, including the secondary and tertiary structures, are stabilized by formation of hydrogen bonds. Through a series of small conformational changes, spatial orientations are modified so as to arrive at the most energetically minimized orientation achievable. The folding of proteins is often facilitated by enzymes known as molecular chaperones. Sterics, bond strain, and angle strain also play major roles in the folding of a protein from its primary sequence to its tertiary structure.

Single tertiary protein structures can also assemble to form protein complexes composed of multiple independently folded subunits. As a whole, this is called a protein's quaternary structure. The quaternary structure is generated by the formation of relatively strong non-covalent interactions, such as hydrogen bonds, between different subunits to generate a functional polymeric enzyme. Some proteins also utilize non-covalent interactions to bind cofactors in the active site during catalysis, however a cofactor can also be covalently attached to an enzyme. Cofactors can be either organic or inorganic molecules which assist in the catalytic mechanism of the active enzyme. The strength with which a cofactor is bound to an enzyme may vary greatly; non-covalently bound cofactors are typically anchored by hydrogen bonds or electrostatic interactions.

Boiling points

Main article: Boiling point

Non-covalent interactions have a significant effect on the boiling point of a liquid. Boiling point is defined as the temperature at which the vapor pressure of a liquid is equal to the pressure surrounding the liquid. More simply, it is the temperature at which a liquid becomes a gas. As one might expect, the stronger the non-covalent interactions present for a substance, the higher its boiling point. For example, consider three compounds of similar chemical composition: sodium n-butoxide (C₄H₉ONa), diethyl ether (C₄H₁₀O), and n-butanol (C₄H₉OH).

Figure 8. Boiling points of 4-carbon compounds

The predominant non-covalent interactions associated with each species in solution are listed in the above figure. As previously discussed, ionic interactions require considerably more energy to break than hydrogen bonds, which in turn are require more energy than dipole–dipole interactions. The trends observed in their boiling points (figure 8) shows exactly the correlation expected, where sodium n-butoxide requires significantly more heat energy (higher temperature) to boil than n-butanol, which boils at a much higher temperature than diethyl ether. The heat energy required for a compound to change from liquid to gas is associated with the energy required to break the intermolecular forces each molecule experiences in its liquid state.

Punishment (psychology)

From Wikipedia, the free encyclopedia

https://en.wikipedia.org/wiki/Punishment_(psychology) 

 

Operant conditioning

Extinction

Reinforcement Increase behavior

Punishment Decrease behavior

Positive reinforcement Add appetitive stimulus following correct behavior

Negative reinforcement

Positive punishment Add noxious stimulus following behavior

Negative punishment Remove appetitive stimulus following behavior

Escape Remove noxious stimulus following correct behavior

Active avoidance Behavior avoids noxious stimulus

In operant conditioning, punishment is any change in a human or animal's surroundings which, occurring after a given behavior or response, reduces the likelihood of that behavior occurring again in the future. As with reinforcement, it is the behavior, not the human/animal, that is punished. Whether a change is or is not punishing is determined by its effect on the rate that the behavior occurs. This is called motivating operations (MO), because they alter the effectiveness of a stimulus. MO can be categorized in abolishing operations, decrease the effectiveness of the stimuli and establishing, increase the effectiveness of the stimuli. For example, a painful stimulus which would act as a punisher for most people may actually reinforce some behaviors of masochistic individuals.

There are two types of punishment, positive and negative. Positive punishment involves the introduction of a stimulus to decrease behavior while negative punishment involves the removal of a stimulus to decrease behavior. While similar to reinforcement, punishment's goal is to decrease behaviors while reinforcement's goal is to increase behaviors. Different kinds of stimuli exist as well. There are rewarding stimuli which are considered pleasant and aversive stimuli, which are considered unpleasant. There are also two types of punishers. There are primary punishers which directly affect the individual such as pain and are a natural response and then there are secondary punishers which are things that are learned to be negative like a buzzing sound when getting an answer wrong on a game show.

Conflicting findings have been found on the effectiveness of the use of punishment. Some have found that punishment can be a useful tool in suppressing behavior while some have found it to have a weak effect on suppressing behavior. Punishment can also lead to lasting negative unintended side effects as well. Punishment has been found to be effective in countries that are wealthy, high in trust, cooperation, and democracy.

Punishment has been used in a lot of different applications. Punishment has been used in applied behavioral analysis, specifically in situations to try and punish dangerous behaviors like head banging. Punishment has also been used to psychologically manipulate individuals to gain control over victims. It has also been used in scenarios where an abuser may try punishment in order to traumatically bond their victim with them. Stuttering therapy has also seen the use of punishment with effective results. Certain punishment techniques have been effective in children with disabilities, such as autism and intellectual disabilities.

Types

See also: Operant conditioning

There are two basic types of punishment in operant conditioning:

• positive punishment, punishment by application, or type I punishment, an experimenter punishes a response by presenting an aversive stimulus into the animal's surroundings (a brief electric shock, for example).
• negative punishment, punishment by removal, or type II punishment, a valued, appetitive stimulus is removed (as in the removal of a feeding dish). As with reinforcement, it is not usually necessary to speak of positive and negative in regard to punishment.

Punishment is not a mirror effect of reinforcement. In experiments with laboratory animals and studies with children, punishment decreases the likelihood of a previously reinforced response only temporarily, and it can produce other "emotional" behavior (wing-flapping in pigeons, for example) and physiological changes (increased heart rate, for example) that have no clear equivalents in reinforcement.

Punishment is considered by some behavioral psychologists to be a "primary process" – a completely independent phenomenon of learning, distinct from reinforcement. Others see it as a category of negative reinforcement, creating a situation in which any punishment-avoiding behavior (even standing still) is reinforced.

Positive

Positive punishment occurs when a response produces a stimulus and that response decreases in probability in the future in similar circumstances.

• Example: A mother yells at a child when they run into the street. If the child stops running into the street, the yelling ceases. The yelling acts as positive punishment because the mother presents (adds) an unpleasant stimulus in the form of yelling.
• Example: A barefoot person walks onto a hot asphalt surface, creating pain, a positive punishment. When the person leaves the asphalt, the pain subsides. The pain acts as positive punishment because it is the addition of an unpleasant stimulus that reduces the future likelihood of the person walking barefoot on a hot surface.

Negative

Negative punishment occurs when a response produces the removal of a stimulus and that response decreases in probability in the future in similar circumstances.

• Example: A teenager comes home after curfew and the parents take away a privilege, such as cell phone usage. If the frequency of the child coming home late decreases, the privilege is gradually restored. The removal of the phone is negative punishment because the parents are taking away a pleasant stimulus (the phone) and motivating the child to return home earlier.
• Example: A child throws a temper tantrum because they want ice cream. Their mother subsequently ignores them, making it less likely the child will throw a temper tantrum in the future when they want something. The removal of attention from his mother is a negative punishment because a pleasant stimulus (attention) is taken away.

Versus reinforcement

Simply put, reinforcers serve to increase behaviors whereas punishers serve to decrease behaviors; thus, positive reinforcers are stimuli that the subject will work to attain, and negative reinforcers are stimuli that the subject will work to be rid of or to end. The table below illustrates the adding and subtracting of stimuli (pleasant or aversive) in relation to reinforcement vs. punishment.

	Rewarding (pleasant) stimulus	Aversive (unpleasant) stimulus
Adding/Presenting	Positive Reinforcement	Positive Punishment
Removing/Taking Away	Negative Punishment	Negative Reinforcement

Types of stimuli and punishers

Rewarding stimuli (pleasant)

A rewarding stimuli is a stimulus that is considered pleasant. For example, a child may be allowed TV time everyday. Punishment often involves the removal of a rewarding stimuli if an undesired action is done. If the child were to misbehave, this rewarding stimulus of TV time would be removed which would result in negative punishment.

Aversive stimuli (unpleasant)

Aversive Stimuli, punisher, and punishing stimulus are somewhat synonymous. Punishment may be used to mean

1. An aversive stimulus
2. The occurrence of any punishing change
3. The part of an experiment in which a particular response is punished.

Some things considered aversive can become reinforcing. In addition, some things that are aversive may not be punishing if accompanying changes are reinforcing. A classic example would be mis-behavior that is 'punished' by a teacher but actually increases over time due to the reinforcing effects of attention on the student.

Primary punishers

Pain, loud noises, foul tastes, bright lights, and exclusion are all things that would pass the "caveman test" as an aversive stimulus, and are therefore primary punishers. Primary punishers can also be loss of money and receiving negative feedback from people.

Secondary punishers

The sound of someone booing, the wrong-answer buzzer on a game show, and a ticket on your car windshield are all things society has learned to think about as negative, and are considered secondary punishers.

Effectiveness

Contrary to suggestions by Skinner and others that punishment typically has weak or impermanent effects, a large body of research has shown that it can have a powerful and lasting effect in suppressing the punished behavior. Furthermore, more severe punishments are more effective, and very severe ones may even produce complete suppression. However, it may also have powerful and lasting side effects. For example, an aversive stimulus used to punish a particular behavior may also elicit a strong emotional response that may suppress unpunished behavior and become associated with situational stimuli through classical conditioning. Such side effects suggest caution and restraint in the use of punishment to modify behavior. Spanking in particular has been found to have lasting side effects. Parents often use spanking to try make their child act better but there is minimal evidence suggesting that spanking is effective in doing so. Some lasting side effects of spanking include lower cognitive ability, lower self-esteem, and more mental health problems for the child. Some side effects can reach into adulthood as well such as antisocial behavior and support for punishment that involves physical force such as spanking. Punishment is more effective in increasing cooperation in high-trust societies than low-trust societies. Punishment was also more effective in countries that have stronger norms for cooperation, high in wealth, and countries that are high-democratic rather than low-democratic.

Importance of contingency and contiguity

One variable affecting punishment is contingency, which is defined as the dependency of events. A behavior may be dependent on a stimulus or dependent on a response. The purpose of punishment is to reduce a behavior, and the degree to which punishment is effective in reducing a targeted behavior is dependent on the relationship between the behavior and a punishment. For example, if a rat receives an aversive stimulus, such as a shock each time it presses a lever, then it is clear that contingency occurs between lever pressing and shock. In this case, the punisher (shock) is contingent upon the appearance of the behavior (lever pressing). Punishment is most effective when contingency is present between a behavior and a punisher. A second variable affecting punishment is contiguity, which is the closeness of events in time and/or space. Contiguity is important to reducing behavior because the longer the time interval between an unwanted behavior and a punishing effect, the less effective the punishment will be. One major problem with a time delay between a behavior and a punishment is that other behaviors may present during that time delay. The subject may then associate the punishment given with the unintended behaviors, and thus suppressing those behaviors instead of the targeted behavior. Therefore, immediate punishment is more effective in reducing a targeted behavior than a delayed punishment would be. However, there may ways to improve the effectiveness of delayed punishment, such as providing verbal explanation, reenacting the behavior, increasing punishment intensity, or other methods.

Applications

Applied behavior analysis

Main article: Applied behavior analysis

Punishment is sometimes used for in applied behavior analysis under the most extreme cases, to reduce dangerous behaviors such as head banging or biting exhibited most commonly by children or people with special needs. Punishment is considered one of the ethical challenges to autism treatment, has led to significant controversy, and is one of the major points for professionalizing behavior analysis. Professionalizing behavior analysis through licensure would create a board to ensure that consumers or families had a place to air disputes, and would ensure training in how to use such tactics properly. (see Professional practice of behavior analysis)

Controversy regarding ABA persists in the autism community. A 2017 study found that 46% of people with autism spectrum undergoing ABA appeared to meet the criteria for post-traumatic stress disorder (PTSD), a rate 86% higher than the rate of those who had not undergone ABA (28%). According to the researcher, the rate of apparent PTSD increased after exposure to ABA regardless of the age of the patient. However, the quality of this study has been disputed by other researchers.

Psychological manipulation

Main article: Psychological manipulation

Braiker identified the following ways that manipulators control their victims:

• Positive reinforcement: includes praise, superficial charm, superficial sympathy (crocodile tears), excessive apologizing, money, approval, gifts, attention, facial expressions such as a forced laugh or smile, and public recognition.
• Negative reinforcement: may involve removing one from a negative situation
• Intermittent or partial reinforcement: Partial or intermittent negative reinforcement can create an effective climate of fear and doubt. Partial or intermittent positive reinforcement can encourage the victim to persist - for example in most forms of gambling, the gambler is likely to win now and again but still lose money overall.
• Punishment: includes nagging, yelling, the silent treatment, intimidation, threats, swearing, emotional blackmail, the guilt trip, sulking, crying, and playing the victim.
• Traumatic one-trial learning: using verbal abuse, explosive anger, or other intimidating behavior to establish dominance or superiority; even one incident of such behavior can condition or train victims to avoid upsetting, confronting or contradicting the manipulator.

Traumatic bonding

Main article: Traumatic bonding

Traumatic bonding occurs as the result of ongoing cycles of abuse in which the intermittent reinforcement of reward and punishment creates powerful emotional bonds that are resistant to change.

Punishment used in stuttering therapy

Main article: Stuttering therapy

Early studies in the late 60's to early 70's have shown that punishment via time-out (a form of negative punishment) can reduce the severity of stuttering in patients. Since the punishment in these studies was time-out which resulted in the removal of the permission to speak, speaking itself was seen as reinforcing which thus made the time-out an effective form of punishment. Some research has also shown that it is not the time-out that is considered punishing but rather the fact that the removal of the permission to speak was seen as punishing because it interrupted the individual's speech.

Punishment in children with disabilities

Some studies have found effective punishment techniques concerning children with disabilities, such as autism and intellectual disabilities. The targeted behaviors were self-injurious behaviors such as head banging, motor, stereotypy, aggression, emesis, or breaking the rules. Some techniques that were used are timeout, overcorrection, contingent aversive, response blocking, and response interruption and redirection (RIRD). Most punishment techniques were used alone or combined with other punishment techniques; however, the use of punishment techniques alone was less effective in reducing targeted behaviors. Timeout was used the most even though it was less effective in reducing targeted behaviors; however, contingent aversive was used the least even though it was more effective in reducing targeted behaviors. Using punishment techniques in combination with reinforcement-based interventions was more effective than a punishment technique alone or using multiple punishment techniques.