in this video we'll look at the type of interactions molecules can have with one another these interactions are called intermolecular forces it's important to understand how and why these interactions occur these intermolecular forces cause very clear differences in physical properties such as boiling point and volatility that is the energy needed for molecules to be moved away from one another and undergo a phase change or solubility and conduct AC ity how molecules can dissolve and be dissolved or viscosity a measure of how liquids flow think for example the difference between pouring water and pouring syrup we've learned to describe the nature of the bonding holding the atoms of a particular substance together but the nature of bonding also affects intermolecular forces we'll examine the three types of intermolecular forces and their impact on the properties of mole molecules coent molecules range from purely non-polar as with the datomic elements where both elements are the same and therefore have the same electro negativity to extremely polar as in the case of water where electrons are shared unfairly across the molecule creating an overall dipole examining the boiling points of calent molecules reveals that properties can vary widely even in molecules with similar structures and molecular weights much of this variation is explained through the difference between intra and Inter molecular forces int molecular forces are the strong calent bonds that hold the molecule together while intermolecular forces are much weaker interactions between molecules resulting from the attractions of molecular dipoles these interactions are weak but the effects are significant there are three types of intermolecular forces the the first called lendo dispersion forces can be found in any atom or molecule no matter their polarity as the electrons of an atom or molecule move throughout the electron cloud the density of the electrons can be irregular the constant movement of electrons create temporary dipoles where the area with an increased electron density temporarily becomes partially negative and the area with a decreased electron density temporarily becomes partially positive any temporary or permanent dipole can cause momentary attraction between molecules or atoms we can see the situation where two atoms or molecules would feel no attractive or repulsive forces as their electrons are evenly distributed causing no partial charges or the situation where atoms or molecules feel attractive forces when their opposite partial charges appear next to each other at the same time and finally the situation where atoms or molecules feel repulsive forces as they face each other with the same partial charges London dispersion forces are similar to traffic on a highway where more vehicles on the road can create traffic jams with groups of cars bunched up near one another like cars on a highway as an atom or molecule increases in size the number of electrons present in its electron cloud also increases allowing for larger shifts of electron density and therefore the creation of larger dipoles all molecules and even larger atoms are polarizable meaning they can create instantaneous dipoles but larger molecules with more electrons are more polarizable and experience stronger Lena dispersion forces we can see this in looking at the boiling point data for datomic Hallens for example we see that bromine and iodine have the highest mass and also the most electrons these two diatomic elements have the highest m melting and boiling points and exist as a liquid and solid respectively at standard ambient temperatures and pressures this is in contrast to chlorine and Florine which are the smallest of the halogens by mass and have the fewest electrons these two diatomic elements have much lower melting and boiling points and are gases at standard ambient temperatures and pressures as they're capable of experiencing much weaker London dispersion forces polar molecules with permanent dipoles will experience more constant and permanent attractive interactions with other molecules of the same substance the stronger the dipole or the greater the molecular polarity the greater the intermolecular forces and the variation in physical properties from other molecules for example the increasing boiling points from hexane to hexany are explained by molecular polarity hexane is considered non-polar which is reflected in its very small dipole moment adding the carbon group and hexanol increases the molecule's polarity giving it a larger dipole moment as a result hexanol experiences the greater intermolecular forces of the two and will have the higher boiling point due to the increased interaction between its molecules hydrogen bonds are the last of our three types of intermolecular forces hydrogen bonding in this context is not a bond rather it is the interaction resulting from the large dipoles that are formed when hydrogen is directly bonded to nitrogen oxygen or Florine these three elements are highly electronegative in comparison to hydrogen and will create very polar bonds with large dipoles when bonded directly to hydrogen further as Period 2 elements nitrogen oxygen and Florine have rather small atomic radi and can more closely approach the hydrogen atom of another molecule the closeness of this interaction paired with the strength of the molecular polarity creates an extreme dipole interaction and thus generates the strongest of the intermolecular forces this is the hydrogen bonding interaction although commonplace and take it for granted life depends on the unique capabilities of water the hydrogen bond in water is a result of the attractive interaction between the partially negative charge on the oxygen atom and the partially positive charge on the hydrogen atom in adjacent water molecules this is a strong intermolecular force that plays a crucial role in various biological and chemical processes When comparing water to methane a molecule with a very similar molecular mass the hydrogen bonding that occurs in water makes its boiling point much higher than methanes as its hydrogen bonding interactions require much more energy to overcome and move water molecules away from one another in order to undergo the liquid to gas phase change Beyond it sign significantly elevated boiling point other unique properties of water include its very high surface tension its role as the universal solvent and the ability via capillary action to move from the soil to the leaves at top a tree sometimes hundreds of feet above the ground water is truly amazing and it's all due to hydrogen bonding by comparing more boiling point data we're able to see the effects of all three types of intermolecular forces let's first decipher our graph each colored series represents the elements of a particular group bonded with hydrogen the green series are hydrogen compounds of group 17 the light red series are hydrogen compounds of group 16 the yellow series are group 15 in the blue series are group 14 H2O HF and NH3 are molecules in which hydrogen is directly bonded to either oxygen Florine or nitrogen these molecules will therefore exhibit hydrogen bonding and have very strong molecular interactions as a result H2O HF and NH3 have higher boiling points than any other compound in their respective series with regard to the first molecule in group 14 note that although CH4 contains hydrogen it does not have hydrogen bonding and is considered non-polar with no dipole dipole interactions by only experiencing London dispersion forces CH4 has has a very low boiling point by examining the light red series Beyond water the compounds H2S H2 SE and h2te are all bent structures that are indeed polar with dipole dipole interactions as we move down group 16 the size of the non-hydrogen atom increases and has more electrons present the slowly increasing boiling point from H2S to h2te is due to increased ining London dispersion forces the larger molecule of h2te is more polarizable and has a greater tendency to develop an instantaneous dipole than H2S finally we could see the effect of dipole dipole interactions by looking at the hydrogen compounds across period 3 we have H2S HCL ph3 and si4 no hydrogen bonding is present in any of these compounds but H2S HCL and pH H 3 are all polar molecules that are capable of dipole dipole interactions si4 however is non-polar and is only capable of experiencing London dispersion forces due to their stronger intermolecular forces H2S HCL and ph3 have greater boiling points than sih4 in summary intermolecular forces known commonly as imfs are interactions between the molecules of a substance and are much weaker than the coent bonds that hold a molecule together London dispersion forces are the weakest of the three imfs although it's important to note that dispersion forces can become significant in large atoms or molecules with many electrons all atoms and molecules are capable of LED dispersion forces dipo dipole interactions are stronger than London dispersion forces and in addition to dispersion exist in polar molecules because of their permanent dipoles finally hydrogen bonding is the strongest of the three intermolecular forces thought of as an extreme dipole interaction any molecule with hydrogen bonding is capable of all three types of intermolecular forces although we focused largely on the effect of intermolecular forces on the boiling points of similar molecules in reality intermolecular forces help to explain differences in many physical properties and help make Life as we know it possible