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Intermolecular Forces & Bonding

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Annelise Crabtree

on 31 October 2012

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Transcript of Intermolecular Forces & Bonding

Intermolecular Forces & Bonding k Ionic Bonds Covalent Bonds Metallic Bonds Ionic bonds are between metals and non-metals. They transfer electrons between elements. The electronegativity scale for ionic bonds is from 1.7 and up. Since electrons are not "owned" by any one element, valence electrons are depicted with brackets and charges, unlike the usual Lewis Dot Structure. The structure of these molecules is a crystal lattice, which offers support. The compounds formed by ionic bonds are usually solids that are soluble in water and are conductors of electricity. Ionic bonds have a high boiling point which indicates strong IMF's.
Ex: NaCl, CaCl, MgO, KBr Covalent bonds are between non-metals. Electrons are shared, equally and un-equally between elements. When electrons are shared equally, the bond is non-polar, and has an electronegativity scale from 0.0 to 0.1. When electrons are shared unequally between elements, the bond is then polar, and has an electronegativity scale from 0.1 to 1.7. When something is polar, (or dipole), there is different charges on opposite ends of a molecule. Covalent Bonds are true molecules, and are usually solids or liquids. They have low melting points, and are usually not soluble in water. Compounds made by covalent bonds are not conductors of electricity and are usually odorous. Metallic Bonds are between metals. Unlike covalent bonds, electrons do not "belong" to any one element. They are part of an "Electron Sea", and the electrons are constantly bouncing from one element to another and do not stay in the same place. Metallic Bonds have very high melting points, and this indicates that the IMF's within these bonds are strong. Compounds created by metallic bonds are solids and are not soluble in water. They are always conductors of electricity, and are malleable, ductile and lustrous in quality.
Ex: CuZn, Au, K, CaCl2 It is the attraction between valence electrons that allows atoms to come together and create a bond, whether it be ionic, covalent or metallic. Electrostatic Attraction Intermolecular Forces Intermolecular Forces Ionic Ionic bonds are the strongest of IMF's. They are the strongest because valence electrons are not shared between elements, and are the result of attraction between oppositely charged ions. These attractions allows the formation of a crystal lattice which is structurally sound and very hard to break. As a result, ionic bonds have a very high melting point, and a low gas pressure, since ionic compounds are less likely to be in gas form. Intermolecular Forces Hydrogen Bonding Occurs when a hydrogen molecule directly bonds to an fluoride, oxygen or nitrogen molecule, (and is a stronger version of a dipole-dipole force). Hydrogen Bonding is the second strongest IMF, or the strongest when concerning just dipoles. This IMF is the attraction between a partially positive hydrogen molecule and the partially negative charge of fluoride, oxygen and or a nitrogen molecule. (FON) Hydrogen Bonds are very difficult to break. Because of this, hydrogen bonds have a high boiling point and a lower gas pressure. These bonds are also extremely polar since it is a stronger version of a
dipole-dipole force.
Ex: H2O, OH, HF Intermolecular Forces Dipole-Dipole Dipole-dipole forces are the third-strongest of IMF's. They are made between polar molecules. Unlike with Van der Waals, the molecules involved with dipole-dipole have permanent dipoles. So, the partially positive end of a molecule is attracted to the partially negative end of a molecule, making this IMF polar. When comparing dipole-dipoles, they are always ranked by the strongest electronegativity. They are not as strong as Hydrogen Bonding or Ionic IMF's, and are not very hard to break. Intermolecular Forces Van der Waals Every bond and compound contains Van der Waals, which is the attraction between temporary dipoles. Since electrons are always moving, sometimes they all come to one side of a molecule, creating partial charges on each end. These temporary dipoles then attract other temporary dipoles, and that there is Van der Waals in action. Ex: H2, H2O, HCl, HF, Si2, SeO Polar "Electron Sea" Crystal Lattice Partial Charges Determining Polarity There are two ways of determining if a compound is polar; looking at its electronegativity and its molecular shape. If the electronegativity is from 0.1-1.7, then it is polar. Also, if its molecular shape is not symmetrical, it is polar. Another way of determining polarity is seeing if any element is "stronger" than another. For example: CH3Cl. Cl is stronger than H, so it will pull the C up, as indicated by the red arrow. Therefore, this compound is polar. This line can act as a mirror; so if this compound was folded in half, the sides would not match up. Since its not symmetrical, it is polar. Electrons are just bouncing around Is just the attraction of opposites Partial positive Partial Negative Ionic Bonds vs. Covalent Bonds In ionic bonding, valence electrons are transferred. But in covalent bonds, electrons are shared. Since electrons are being shared with covalent bonds, covalent compounds can be shown with the traditional Lewis Dot Diagrams. Since electrons are being transferred in ionic bonding, these compounds are shown with brackets and charges. Also, covalent bonds are between two nonmetals, while ionic bonds are between a metal and a nonmetal. Ionic bonds are stronger than covalent bonds because ionic bonds are usually solids. Of the three states of matter, solids are the most sturdy and durable due to strong IMF's. So, since ionic bonds are mostly solids and have strong IMF's, they are naturally more strong compared to covalent bonds which are usually liquids and solids. Liquids are not as weak as gases, but not as strong as solids. IMF's are weakened in liquids, causing them to flow more easily. (Gases have no IMF's, allowing the molecules to bounce around in the air). So, covalent bonds are weaker than ionic bonds, and therefore easier to break. Ionic/Covalent vs. Metallic Bonds Unlike either ionic bonding or covalent bonding, in metallic bonding, electrons do not belong to any element, or are shared between elements. Electrons are in an "Electron Sea", and constantly float around the molecules. Metallic compounds are between two metals, and not two nonmetals or a nonmetal and a metal. They are always solids, and not liquids (covalent). They have this in common with ionic compounds, and some covalent compounds. Metallic bonds are the strongest of all three bonds because of their lattice structure; quite like ionic bonds. But metallic lattices are stronger because atoms are more closely packed together than in ionic bonds. And in covalent compounds' case, their form comes in true molecules. So metallic bonds are stronger than ionic and covalent bonds. More Ionic Bonds vs. Covalent Bonds Ionic compounds dissolve is water, while covalent compounds usually do not. Ionic compounds, since they have stronger IMF's, have a higher boiling point than the weaker IMF's of covalent compounds. Ionic compounds are conductors of electricity while covalent compounds are not. Ionic bonds have a crystal lattice structure while covalent bonds are in the form of true molecules. More Ionic/Covalent vs. Metallic Bonding Compared to ionic and covalent compounds from the last circle, metallic compounds (like ionic) have a lattice structure. They are not soluble in water. They are conductors of electricity, and are malleable, ductile and lustrous in quality. Isotopes Are molecules of an element that have a different number of neutrons. But, they still behave chemically the same as molecules that are not isotopes of a certain element. So in other words, they bond the same as non-isotopes; ionic, covalent and metallic. First off, elements bond to form a perfect octet, or to gain eight valence electrons. They do this in order to become more stable. The only elements that have a perfect octet are the Noble Gases.
There are three types of bonds:
* Ionic Bonds
* Covalent Bonds
* Metallic Bonds Electron Configuration Electron configuration shows how many valence electrons an element has. All chemical bonds are made in order to form an octet. This directly applies to bonding because it tells us how an element will react chemically. For Example, Oxygen (O): 1S2, 2S, 2P4, and has 6 valence electrons. So in order to form an octet, it will need to gain 2 more electrons.
Works Cited

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