The strength of these attractions also determines what changes in temperature and pressure are needed to effect a phase transition. They arise from the formation of temporary, instantaneous polarities across a molecule from circulations of electrons. Although the lone pairs in the chloride ion are at the 3-level and would not normally be active enough to form hydrogen bonds, in this case they are made more attractive by the full negative charge on the chlorine. Fully explain how you determined this. Although CH bonds are polar, they are only minimally polar. intermolecular forces (check all that apply) compound dispersion dipole hydrogen-bonding carbon monoxide Cl2 chlorine HBrO hypobromous acid NOC nitrosyl chloride . See Answer . Determine the intermolecular forces in the compounds and then arrange the compounds according to the strength of those forces. Intermolecular forces are electrostatic in nature; that is, they arise from the interaction between positively and negatively charged species. Similarly, solids melt when the molecules acquire enough thermal energy to overcome the intermolecular forces that lock them into place in the solid. Chang, Raymond. Table \(\PageIndex{1}\) lists the exponents for the types of interactions we will describe in this lesson. Covalent bonds with these elements are very polar, resulting in a partial negative charge () on the O, N, or F. This partial negative charge can be attracted to the partial positive charge (+) of the hydrogen in an XH bond on an adjacent molecule. The interaction between two molecules can be decomposed into different combinations of moment-moment interactions. Chemical bonds (e.g., covalent bonding) are intramolecular forces which hold atoms together as molecules. The hydrogen bonding is limited by the fact that there is only one hydrogen in each ethanol molecule with sufficient + charge. Within a vessel, water molecules hydrogen bond not only to each other, but also to the cellulose chain which comprises the wall of plant cells. Hydrogen bonding can occur between ethanol molecules, although not as effectively as in water. The first two are often described collectively as van der Waals forces. For example: monopole-monopole is a charge-charge interaction (Equation \(\ref{Col}\)), monopole-dipole, dipole-dipole, charge-quadrupole, dipole-quadrupole, quadrupole-quadrupole, charge-octupule, dipole-octupole, quadrupole-octupole, octupole-octople etc. a. Ion-dipole forces explanations are helpful! It is important to realize that hydrogen bonding exists in addition to van der Waals attractions. In contrast, each oxygen atom is bonded to two H atoms at the shorter distance and two at the longer distance, corresponding to two OH covalent bonds and two OH hydrogen bonds from adjacent water molecules, respectively. Thus far we have considered only interactions between polar molecules, but other factors must be considered to explain why many nonpolar molecules, such as bromine, benzene, and hexane, are liquids at room temperature, and others, such as iodine and naphthalene, are solids. The boiling points of ethanol and methoxymethane show the dramatic effect that the hydrogen bonding has on the stickiness of the ethanol molecules: The hydrogen bonding in the ethanol has lifted its boiling point about 100C. There are no hydrogen bonds, because NF3 doesn't have any HF , HO , or HN bonds. GeCl4 (87C) > SiCl4 (57.6C) > GeH4 (88.5C) > SiH4 (111.8C) > CH4 (161C). These interactions become important for gases only at very high pressures, where they are responsible for the observed deviations from the ideal gas law at high pressures. As shown in part (a) in Figure \(\PageIndex{3}\), the instantaneous dipole moment on one atom can interact with the electrons in an adjacent atom, pulling them toward the positive end of the instantaneous dipole or repelling them from the negative end. The reason for this trend is that the strength of London dispersion forces is related to the ease with which the electron distribution in a given atom can be perturbed. Consequently, even though their molecular masses are similar to that of water, their boiling points are significantly lower than the boiling point of water, which forms four hydrogen bonds at a time. Examples include permanent monopole (charge) - induced dipole interaction, permanent dipole - induced dipole interaction, permanent quadrupole-induced dipole interaction etc. This prevents the hydrogen bonding from acquiring the partial positive charge needed to hydrogen bond with the lone electron pair in another molecule. Hydrogen bonds in HF(s) and H2O(s) (shown on the next page) are intermediate in strength within this range. Molecules with higher molecular weights have more electrons, which are generally more loosely held. (Forces that exist within molecules, such as chemical bonds, are called intramolecular forces.) The hydrogen bonding is limited by the fact that there is only one hydrogen in each ethanol molecule with sufficient, lone pairs on the oxygen are still there, but the. compound intermolecular forces (check all that apply) dispersion dipole hydrogen-bonding SiH silane . The bridging hydrogen atoms are not equidistant from the two oxygen atoms they connect, however. Neopentane is almost spherical, with a small surface area for intermolecular interactions, whereas n-pentane has an extended conformation that enables it to come into close contact with other n-pentane molecules. intermolecular forces (check all that apply) compound dispersion dipole hydrogen-bonding carbon monoxide hypobromous acid nitrogen tribromide chlorine This problem has been solved! (see Polarizability). Transcribed Image Text: intermolecular forces compound (check all that apply) dispersion dipole hydrogen-bonding hydrogen chloride hydrogen fluoride carbon dioxide nitrogen tribromide This occurs when two functional groups of a molecule can form hydrogen bonds with each other. Identify the intermolecular forces in each compound and then arrange the compounds according to the strength of those forces. 2.10: Intermolecular Forces (IMFs) - Review is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts. Methane (CH) London dispersion forces . Consider a pair of adjacent He atoms, for example. This is the expected trend in nonpolar molecules, for which London dispersion forces are the exclusive intermolecular forces. Arrange ethyl methyl ether (CH3OCH2CH3), 2-methylpropane [isobutane, (CH3)2CHCH3], and acetone (CH3COCH3) in order of increasing boiling points. Intermolecular forces (IMF) can be qualitatively ranked using Coulomb's Law: Less than 0.40. Based on the IMF present in each of the molecules below, predict the relative boiling points of each of the substances below. Identify the strongest intermolecular force present in pure samples of the following substances: Identify the strongest intermolecular force operating in the condensed phases of the following substances. The four compounds are alkanes and nonpolar, so London dispersion forces are the only important intermolecular forces. However, the relevant moments that is important for the IMF of a specific molecule depend uniquely on that molecules properties. Although hydrogen bonds are significantly weaker than covalent bonds, with typical dissociation energies of only 1525 kJ/mol, they have a significant influence on the physical properties of a compound. On average, the two electrons in each He atom are uniformly distributed around the nucleus. London dispersion forces and dipole-dipole forces are collectively known as van der Waals forces. Arrange n-butane, propane, 2-methylpropane [isobutene, (CH3)2CHCH3], and n-pentane in order of increasing boiling points. As we have seen, the model of an ideal gas assumes that the gas particles (molecules or atoms) have virtually no forces of attraction between them, are widely separated, and are constantly moving with high velocity and kinetic energy. Intermolecular forces are generally much weaker than covalent bonds. This effect, illustrated for two H2 molecules in part (b) in Figure \(\PageIndex{3}\), tends to become more pronounced as atomic and molecular masses increase (Table \(\PageIndex{2}\)). Except in some rather unusual cases, the hydrogen atom has to be attached directly to the very electronegative element for hydrogen bonding to occur. Consequently, N2O should have a higher boiling point. Often, but not always, these interactions can be ranked in terms of strengths with of interactions involving lower number of moments dominating those with higher moments. The three compounds have essentially the same molar mass (5860 g/mol), so we must look at differences in polarity to predict the strength of the intermolecular dipoledipole interactions and thus the boiling points of the compounds. For example, Xe boils at 108.1C, whereas He boils at 269C. For example, all the following molecules contain the same number of electrons, and the first two are much the same length. Imagine the implications for life on Earth if water boiled at 130C rather than 100C. Legal. This, without taking hydrogen bonds into account, is due to greater dispersion forces (see Interactions Between Nonpolar Molecules). If you are interested in the bonding in hydrated positive ions, you could follow this link to co-ordinate (dative covalent) bonding. Water (HO) hydrogen bonding . The strength of the induced dipole moment, \(\mu_{induced}\), is directly proportional to the strength of the electric field, \(E\) of the permanent moment with a proportionality constant \(\alpha\) called the polarizability. Water (H20) dipole-dipole. Such molecules will always have higher boiling points than similarly sized molecules which don't have an -O-H or an -N-H group. Within a vessel, water molecules hydrogen bond not only to each other, but also to the cellulose chain which comprises the wall of plant cells. Thus London dispersion forces are responsible for the general trend toward higher boiling points with increased molecular mass and greater surface area in a homologous series of compounds, such as the alkanes (part (a) in Figure \(\PageIndex{4}\)). It can be classified as ionic force, dipole-dipole force, H-bonding, or London dispersion force depending on how the electrons are distributed around the substance's particle. 3) silicon tetrafluoride (SiF4) London dispersion forces 4) nitrogen tribromide (NBr3) dipole-dipole forces 5) water (H2O) hydrogen bonding 6) methane (CH4) London dispersion forces7) benzene (C6H6) London dispersion forces 8) ammonia (NH3) ) hydrogen bonding 9) methanol (CH3OH))hydrogen bonding This process is called hydration. Recall that the attractive energy between two ions is proportional to 1/r, where r is the distance between the ions. The strength of the electric field causes the distortion in the molecule. Changing those conditions can induce a change in the state of the substance, called a phase transition. The hydrogen atom is then left with a partial positive charge, creating a dipole-dipole attraction between the hydrogen atom bonded to the donor, and the lone electron pair on the, hydrogen bonding occurs in ethylene glycol (C, The same effect that is seen on boiling point as a result of hydrogen bonding can also be observed in the, Hydrogen bonding plays a crucial role in many biological processes and can account for many natural phenomena such as the, The cohesion-adhesion theory of transport in vascular plants uses hydrogen bonding to explain many key components of water movement through the plant's xylem and other vessels. You should try to answer the questions without accessing the Internet. Because all molecules have electrons, all molecular substances have London dispersion forces, regardless of whether they are polar or non-polar. 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. Decide which intermolecular forces act between the molecules of each compound in the table below. If ice were denser than the liquid, the ice formed at the surface in cold weather would sink as fast as it formed. These interactions occur because of hydrogen bonding between water molecules around the hydrophobe and further reinforce conformation. Nitrogen tribromide | Br3N - PubChem Apologies, we are having some trouble retrieving data from our servers. a. Instantaneous dipoleinduced dipole interactions between nonpolar molecules can produce intermolecular attractions just as they produce interatomic attractions in monatomic substances like Xe. As we have seen, the model of an ideal gas assumes that the gas particles (molecules or atoms) have virtually no forces of attraction between them, are widely separated, and are constantly moving with high velocity and kinetic energy. The former is termed an intramolecular attraction while the latter is termed an intermolecular attraction. Nonpolar covalent difference in electronegativity. Because each water molecule contains two hydrogen atoms and two lone pairs, a tetrahedral arrangement maximizes the number of hydrogen bonds that can be formed. In the case of liquids, molecular attractions give rise to viscosity, a resistance to flow. In fact, the ice forms a protective surface layer that insulates the rest of the water, allowing fish and other organisms to survive in the lower levels of a frozen lake or sea. Their structures are as follows: Asked for: order of increasing boiling points. Interactions between these temporary dipoles cause atoms to be attracted to one another. PUGVIEW FETCH ERROR: 403 Forbidden National Center for Biotechnology Information 8600 Rockville Pike, Bethesda, MD, 20894 USA Contact Policies FOIA HHS Vulnerability Disclosure National Library of Medicine National Institutes of Health In this section, we explicitly consider three kinds of intermolecular interactions: There are two additional types of electrostatic interaction that you are already familiar with: the ionion interactions that are responsible for ionic bonding and the iondipole interactions that occur when ionic substances dissolve in a polar substance such as water. On average, however, the attractive interactions dominate. These additional forces of attraction must be overcome in a transition to a less-ordered phase (e.g., solid to liquid, liquid to gas), so substances with dipole-dipole attractions between their molecules tend to have higher melting points and boiling points than comparable compounds composed of nonpolar molecules, which only have London dispersion intermolecular forces. Silicon Tetrafluoride (SiF) London dispersion forces. The van, attractions (both dispersion forces and dipole-dipole attractions) in each will be much the same. Thus, we see molecules such as PH3, which no not partake in hydrogen bonding. The substance with the weakest forces will have the lowest boiling point. The predicted order is thus as follows, with actual boiling points in parentheses: He (269C) < Ar (185.7C) < N2O (88.5C) < C60 (>280C) < NaCl (1465C). Polar covalent bonds behave as if the bonded atoms have localized fractional charges that are equal but opposite (i.e., the two bonded atoms generate a dipole). Although the mix of types and strengths of intermolecular forces determines the state of a substance under certain conditions, in general most substances can be found in any of the three states under appropriate conditions of temperature and pressure. For example, intramolecular hydrogen bonding occurs in ethylene glycol (C2H4(OH)2) between its two hydroxyl groups due to the molecular geometry. KBr (1435C) > 2,4-dimethylheptane (132.9C) > CS2 (46.6C) > Cl2 (34.6C) > Ne (246C). Top. As a result, the CO bond dipoles partially reinforce one another and generate a significant dipole moment that should give a moderately high boiling point. Because a hydrogen atom is so small, these dipoles can also approach one another more closely than most other dipoles. In addition to being present in water, hydrogen bonding is also important in the water transport system of plants, secondary and tertiary protein structure, and DNA base pairing. Methanol (CH3OH) hydrogen bonding. Similarly, if a molecule does not have a dipole moment nor monopole moment, then quadrupolar interactions will be important. In order for this to happen, both a hydrogen donor an acceptor must be present within one molecule, and they must be within close proximity of each other in the molecule. Because ice is less dense than liquid water, rivers, lakes, and oceans freeze from the top down. Because the electron distribution is more easily perturbed in large, heavy species than in small, light species, we say that heavier substances tend to be much more polarizable than lighter ones. For the most part, only compounds in which hydrogen is covalently bonded to O, N, or F are candidates for hydrogen bonding. to large molecules like proteins and DNA. PH3 exhibits a trigonal pyramidal molecular geometry like that of ammmonia, but unlike NH3 it cannot hydrogen bond. Indicate which of the following properties will increase, decrease or remain unaffected by an increase in the strength of the intermolecular forces? The secondary structure of a protein involves interactions (mainly hydrogen bonds) between neighboring polypeptide backbones which contain Nitrogen-Hydrogen bonded pairs and oxygen atoms. This question was answered by Fritz London (19001954), a German physicist who later worked in the United States. Water frequently attaches to positive ions by co-ordinate (dative covalent) bonds. Since SiF4 has a greater molecular mass than SiH4, therefore SiF4 has a greater London dispersion force and a greater boiling point. The most significant intermolecular force for this substance would be dispersion forces. Nitrogen is a chemical element with the atomic number 7 and the symbol N. Two atoms of the element bind to form N2, a colourless and odourless diatomic gas, at standard temperature and pressure. However, when we consider the table below, we see that this is not always the case. Hydrogen bonding plays a crucial role in many biological processes and can account for many natural phenomena such as the Unusual properties of Water. Weakest intermolecular force. In contrast to intra molecular forces, such as the covalent bonds that hold atoms together in molecules and polyatomic ions, inter molecular forces hold molecules . The diagram shows the potential hydrogen bonds formed to a chloride ion, Cl-. Hydrogen bonding also occurs in organic molecules containing N-H groups - in the same sort of way that it occurs in ammonia. Identify the most significant intermolecular force in each substance. Because the electrons are in constant motion, however, their distribution in one atom is likely to be asymmetrical at any given instant, resulting in an instantaneous dipole moment. These result in much higher boiling points than are observed for substances in which London dispersion forces dominate, as illustrated for the covalent hydrides of elements of groups 1417 in Figure \(\PageIndex{5}\). Also, the absence of intermolecular forces above the surface of a liquid results in surface tension, the development of a skin on the surface, which causes beading of liquid droplets and also allows light objects to rest on a liquid surface without sinking (e.g., water bugs). The LibreTexts libraries arePowered by NICE CXone Expertand are supported by the Department of Education Open Textbook Pilot Project, the UC Davis Office of the Provost, the UC Davis Library, the California State University Affordable Learning Solutions Program, and Merlot. Intermolecular hydrogen bonds occur between separate molecules in a substance. Polar Covalent difference in electronegativity. With stronger intermolecular forces or lower kinetic energy, those forces may draw molecules closer together, resulting in a condensed phase. Solids have stronger intermolecular forces, making them rigid, with essentially no tendency to flow. What is the predominant intermolecular force in ? 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 are not the same. We can examine which of these forces apply to tetrabromomethane (carbon tetrabromide). Most substances can exist in either gas, liquid, or solid phase under appropriate conditions of temperature and pressure. 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. In truth, there are forces of attraction between the particles, but in a gas the kinetic energy is so high that these cannot effectively bring the particles together. Give an explanation in terms of IMF for the following differences in boiling point. Since the hydrogen donor is strongly electronegative, it pulls the covalently bonded electron pair closer to its nucleus, and away from the hydrogen atom. Hydrogen bond strengths typically are in the range 4 - 46 kJ/mol, much less than the strengths of typical covalent bonds. When any molecules are in direct contact a strong repulsion force kicks in. The hydrogen atom is then left with a partial positive charge, creating a dipole-dipole attraction between the hydrogen atom bonded to the donor, and the lone electron pair on the accepton. Intermolecular Forces and Interactions (Worksheet) is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts. London was able to show with quantum mechanics that the attractive energy between molecules due to temporary dipoleinduced dipole interactions falls off as 1/r6. Intermolecular Forces of Attraction: The intermolecular force of attraction, usually abbreviated as IMFA, is the force that keeps the particles of a substance together. The first two interactions are the most relevant for our discussion. Boron trifluoride (BF3) Dispersion forces Boron difluoride (BF2H) Dipole forces Hydrogen (H2) london forces Carbon Monoxide (CO) london forces Silicon Tetrafluoride (SiF4) london forces Nitrogen Tribromide (NBr3) dipole-dipole; london forces Students also viewed Intermolecular forces 24 terms Joel_Varner6 Intermolecular Forces 18 terms Nitrogen tribromide(NBr) dipole dipole forces. The greater the strength of the intermolecular forces, the more likely the substance is to be found in a condensed state; i.e., either a liquid or solid. Between ~0.41 to ~2.0. \(A\) and \(B\) are proportionality constants and \(n\) and \(m\) are integers. determine the dominant intermolecular forces (IMFs) of organic compounds. Sets with similar terms. Any molecule which has a hydrogen atom attached directly to an oxygen or a nitrogen is capable of hydrogen bonding. This process is called, If you are interested in the bonding in hydrated positive ions, you could follow this link to, They have the same number of electrons, and a similar length to the molecule. Ethyl methyl ether has a structure similar to H2O; it contains two polar CO single bonds oriented at about a 109 angle to each other, in addition to relatively nonpolar CH bonds. The following data for the diatomic halogens nicely illustrate these trends. Nitrogen tribromide is slightly polar in nature. Methane and its heavier congeners in group 14 form a series whose boiling points increase smoothly with increasing molar mass. The polar covalent bond is much stronger in strength than the dipole-dipole interaction. These interactions occur because of hydrogen bonding between water molecules around the, determine the dominant intermolecular forces (IMFs) of organic compounds. a covalent bond in which the electrons are shared equally by the two atoms. Correspondingly, if \(q_1\) and \(q_2\) have the same sign, then the force is negative (i.e., a repulsive interaction). The hydrogen bonding IMF is a special moment-moment interaction between polar groups when a hydrogen (H) atom covalently bound to a highly electronegative atom such as nitrogen (N), oxygen (O), or fluorine (F) experiences the electrostatic field of another highly electronegative atom nearby. Of the two butane isomers, 2-methylpropane is more compact, and n-butane has the more extended shape. These bonds are broken when the compound undergoes a phase change. When an ionic substance dissolves in water, water molecules cluster around the separated ions. Acetone contains a polar C=O double bond oriented at about 120 to two methyl groups with nonpolar CH bonds. They have the same number of electrons, and a similar length to the molecule. Doubling the distance (r 2r) decreases the attractive energy by one-half. This results in a hydrogen bond. Since the hydrogen donor is strongly electronegative, it pulls the covalently bonded electron pair closer to its nucleus, and away from the hydrogen atom. Intermolecular Forces: Intermolecular forces refer to the bonds that occur between molecules. N2 constitutes approximately 78 % of the Earth's atmosphere, making it the most abundant uncombined element. The strengths of London dispersion forces also depend significantly on molecular shape because shape determines how much of one molecule can interact with its neighboring molecules at any given time. The size of donors and acceptors can also effect the ability to hydrogen bond. Those substances which are capable of forming hydrogen bonds tend to have a higher viscosity than those that do not. We see that H2O, HF, and NH3 each have higher boiling points than the same compound formed between hydrogen and the next element moving down its respective group, indicating that the former have greater intermolecular forces. So now we can define the two forces: Intramolecular forces are the forces that hold atoms together within a molecule. The most significant force in this substance is dipole-dipole interaction. Chemistry . It bonds to negative ions using hydrogen bonds. Boiling point increases due to the increasing molar masses, increasing surface tension, increasing intermolecular forces. Solving this integral is beyond the scope of Chem 2BH, but the gist is important: Dipole-dipole forces of attraction exist between molecules that are polar those that have a permanent dipole moment. Transcribed Image Text: Decide which intermolecular forces act between the molecules of each compound in the table below. is due to the additional hydrogen bonding. Furthermore, \(H_2O\) has a smaller molar mass than HF but partakes in more hydrogen bonds per molecule, so its boiling point is consequently higher. In larger atoms such as Xe, however, the outer electrons are much less strongly attracted to the nucleus because of filled intervening shells. To predict the relative boiling points of the other compounds, we must consider their polarity (for dipoledipole interactions), their ability to form hydrogen bonds, and their molar mass (for London dispersion forces). where \(q_1\) and \(q_2\) are charges and \(r\) is the distance between them. The net effect is that the first atom causes the temporary formation of a dipole, called an induced dipole, in the second. What kind of attractive forces can exist between nonpolar molecules or atoms? This can account for the relatively low ability of Cl to form hydrogen bonds. Hydrogen bonding is present abundantly in the secondary structure of proteins, and also sparingly in tertiary conformation. All molecules, whether polar or nonpolar, are attracted to one another by London dispersion forces in addition to any other attractive forces that may be present. A Of the species listed, xenon (Xe), ethane (C2H6), and trimethylamine [(CH3)3N] do not contain a hydrogen atom attached to O, N, or F; hence they cannot act as hydrogen bond donors. A general empirical expression for the potential energy between two particles can be written as, \[V(r) = Ar^{-n} + Br^{-m} \label{7.2.1} \]. There are multiple "flavors" of IMF, but they originate from Equation \(\ref{Col}\), but differ in terms of charge distributions. General Chemistry:The Essential Concepts. Sketch the orientations of molecules and/or ions involved in the following intermolecular attractive forces. Nonetheless, hydrogen bond strength is significantly greater than either London dispersion forces or dipole-dipole forces. ionic. The properties of liquids are intermediate between those of gases and solids but are more similar to solids. In 1930, London proposed that temporary fluctuations in the electron distributions within atoms and nonpolar molecules could result in the formation of short-lived instantaneous dipole moments, which produce attractive forces called London dispersion forces between otherwise nonpolar substances. Draw the hydrogen-bonded structures. The boiling point of the, 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. The \(B\) coefficient is negative for attractive forces, but it will become positive for electrostatic repulsion between like charges. The LibreTexts libraries arePowered by NICE CXone Expertand are supported by the Department of Education Open Textbook Pilot Project, the UC Davis Office of the Provost, the UC Davis Library, the California State University Affordable Learning Solutions Program, and Merlot. In the structure of ice, each oxygen atom is surrounded by a distorted tetrahedron of hydrogen atoms that form bridges to the oxygen atoms of adjacent water molecules. In addition, the attractive interaction between dipoles falls off much more rapidly with increasing distance than do the ionion interactions. There are 3 main types of intermolecular forces between molecules: hydrogen bonding, dipole-dipole, and London dispersion forces. Compounds such as HF can form only two hydrogen bonds at a time as can, on average, pure liquid NH3. CHEM-Intermolecular Forces Mastering Chemistry. You'll get a detailed solution from a subject matter expert that helps you learn core concepts.
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