Notes: Covalent Compounds

Transcription

1 Notes: Covalent Compounds There are two ways that elements want to be like the nearest noble gas: 1) Gain or lose electrons to form an ionic compound. 2) Share electrons with other elements to form covalent compounds. How/why does this happen? Whenever two nonmetals bond to each other, electrons don t get transferred because both elements have similar electronegativities (i.e. they both want to gain electrons to be like the nearest noble gas). o In ionic bonding (as with NaCl, Cl is electronegative and Na isn t, so Na doesn t mind giving electrons to Cl). This way, both fill their octets. o In covalent bonding (as with F 2 ), both elements have similar electronegativities so neither will give electrons to the other. As a result, they re forced to share electrons.

2 How to two nonmetals share electrons? Let s look at what happens when two fluorine atoms bond: 7 valence e - 7 valence e - If they stick together so they share electrons, they ll both be convinced they have 8 electrons! The s- and p- electrons that they have since the last noble gas are called valence electrons. Because nonmetals bond by sharing valence electrons, they re called covalent compounds. Definitions: Covalent compound: A compound formed when nonmetals bond by sharing two or more valence electrons. Covalent bond: A chemical bond formed when nonmetal atoms share two valence electrons. More examples: oxygen bonds with oxygen: o How many valence electrons does each one have? (6) o Arrange the valence electrons to show them as unpaired as is possible, the same way we do for orbital filling diagrams. Two pairs of valence electrons are shared, forming an O=O double bond o Because double bonds involve sharing 4 electrons, double bonds are much stronger than single bonds! This energy is called the bond dissociation energy.

3 More examples of covalent bonding: o N 2 (explain how triple bonds have 6 shared electrons so are stronger have higher bond dissociation energies - than double or single bonds) o F 2 O, and so forth. Properties of covalent compounds: All properties of covalent compounds are determined by the fact that covalent compounds form molecules, while ionic compounds form crystals. A good model for thinking of this: o Ionic compounds are like stacks of Legos all locked together into a big block. o Covalent compounds are like rubber balls thrown together into the same bucket. 1) Covalent compounds have low melting and boiling points: In ionic compounds, all of the atoms are magnetically stuck together in great big crystals To melt an ionic compound, you need to overcome the magnetic force for all the ions in the crystal. In covalent compounds, the molecules only have very weak forces (called Van der Waals forces) holding them to each other. As a result, covalent compounds can be found as solids, liquids, or gases at standard room conditions. To melt a covalent compound, you DO NOT BREAK COVALENT BONDS!!!! All you do is separate the molecules from each other. Since the Van der Waals forces holding them next to each other are very weak, it doesn t require much energy to do this.

4 2) Covalent compounds are soft and squishy Explain how the rubber ball vs. Lego block analogy works here. Also explain that this is only a general property some covalent compounds can be quite hard. 3) Covalent compounds usually don t dissolve in water as well as ionic compounds. Explain how water pulls salts apart and show briefly how the water molecules have less success grabbing onto covalent compounds. 4) Covalent compounds don t conduct electricity (either when solids, melted, or dissolved in water) Since there are no charged particles (as in ionic compounds) and no delocalized electrons (as in metals), they don t conduct electricity at all. 5) Covalent compounds sometimes burn. Compounds that burn usually contain carbon and hydrogen. Because carbon and hydrogen form covalent molecules when they bond with each other, some covalent compounds are able to burn. Covalent compounds that don t contain carbon and hydrogen usually don t burn.

5 Notes: Naming Covalent Compounds Writing the names of covalent compounds All names have two words: The first word is the same as the name of the first element in the formula. The second word is the same as the name of the second element with -ide at the end. Use prefixes in front of each word to indicate how many of each atom are present in the compound. Number of atoms Prefix 1 mono- (only for oxygen) 2 di- 3 triand so on until 10 and so on until deca- Exceptions to these rules: o Because some compounds were discovered before the elements were known, some covalent compounds have names that don t match their formulas. Three big examples: Water H 2 O Ammonia NH 3 CH 4 methane o Elements: If a compound contains only one element, the name of the compound is the same as the name of the element. F 2 fluorine P 4 phosphorus

6 Examples: o PCl 3 = phosphorus tribromide o S 8 = sulfur o N 2 S 3 = dinitrogen trisulfide o CO = carbon monoxide Writing formulas if given the names: Explain how to do this for most covalent compounds. Explain the element formulas: o Big 7 and the weirdo are diatomic. o Phosphorus is P 4 o Sulfur is S 8 o Everything else has the formula of the atomic symbol. Some examples. Naming acids (compounds that start with H): For acids without oxygen, the name is written as hydro[anion]ic acid. o H 2 S = hydrosulfuric acid o HBr = hydrobromic acid o HCN = hydrocyanic acid For acids that contain oxygen (called oxyacids ): The name of the acid is [anion name][suffix] acid. o The suffix depends on the name of the anion: If the anion ends in -ate, the suffix is -ic. HNO 3 = nitric acid H 2 SO 4 = sulfuric acid H 3 PO 4 = phosphoric acid If the anion ends in -ite, the suffix is ous. HNO 2 = nitrous acid H 2 SO 3 = sulfurous acid

7 Types of formulas: There are three types of formulas that are used to describe chemical compounds: Molecular formulas: Tell you how many of each type of atom are present in a molecule o These are the formulas that you re used to working with. o Example: C 6 H 6 has 6 carbon atoms and 6 hydrogen atoms. Empirical formulas: Reduced molecular formulas that tell you the ratios of the elements to each other. o Explain that historically, this was due to the use of combustion analysis, which gave only ratios of elements to one another. o Example: C 6 H 6 reduces to CH. o Also explain that many compounds may have the same empirical formula, making it impossible to identify a specific compound by an empirical formula alone. After all, the empirical formula CH could be C 2 H 2, C 3 H 3, and so forth Structural formulas: Formulas that not only tell you how many of each type of atom are present, but also tell you where they are. o Basically, these are pictures that show you all of the atoms in a molecule. o You saw these when we first learned about covalent compounds. o There are many types of structural formulas (briefly explain the ones used in biochemistry and stereochemistry), but the type we re primarily going to be using are called Lewis structures.

8 How to draw structural formulas: Lewis Structures Example: CH 4 1) Count the total number of valence electrons in the molecule. To do this, find the total number of s- and p-electrons each atom has since the last noble gas, and add them together. For example, hydrogen has one. If the molecule has a negative charge shown in the formula, add that to the total number of valence electrons. If there is positive charge shown, subtract it from the number of valence electrons. For our example: C: 4 electrons x 1 atom = 4 valence electrons H: 1 electron x 4 atoms = 4 valence electrons Total: 8 valence electrons 2) Find the number of octet electrons for the molecule. The rules for doing this: o Hydrogen ALWAYS wants 2 octet electrons! o Beryllium ALWAYS wants 4 octet electrons! o Boron wants 6 electrons for neutral molecules, 8 if it s in an anion. o ALL OTHER ELEMENTS ALWAYS WANT 8 ELECTRONS! For our example: C: 8 octet electrons x 1 atom = 8 octet electrons H: 2 octet electrons x 4 atoms = 8 octet electrons Total: 16 octet electrons

9 3) Subtract the number of valence electrons from the number of octet electrons to find the number of bonding electrons. Example: 16 8 = 8 bonding electrons 4) Divide the number of bonding electrons by two to find the number of bonds in the molecule. You divide by two because there are two electrons in every covalent bond. Example: 8 / 2 = 4 bonds 5) Draw an arrangement for the atoms that has the number of bonds you found in step 4 and follows these rules: The Rules: o Hydrogen and the halogens ALWAYS bond once! NEVER MORE! o Oxygen s family and beryllium bond twice in neutral molecules, once or twice of the molecule has charge. o Nitrogen s family and boron bond three times in neutral molecules and 2, 3, or 4 times in molecules with charge. o Carbon s family ALWAYS bonds 4 times! o When you can stick all of the atoms together so this works, you re done! Some handy suggestions for sticking molecules together: o If you bond all of the atoms together with one bond and there are bonds left over, then you may need to make double and/or triple bonds. o The molecule that s nearest the left side of the periodic table is probably in the middle of the molecule. o It s not correct (in most cases) to draw structures in which the atoms form a ring. For our example: Show the Lewis structure of methane and explain how it comes together.

10 6) Add pairs of electrons to the structure until all atoms have the number of electrons around them that we said they needed in step 2: In our example, we need to look at the number of electrons around each of the atoms in the Lewis structure. Carbon has eight electrons around it. Because carbon only wants eight electrons, we don t need to add any more lone pairs of electrons. Each hydrogen wants two electrons around it. Because carbon needs only two electrons, we don t need to add any more lone pairs of electrons on hydrogen. o Handy hint: You NEVER need to add any pairs of electrons to carbon s family, the halogens, or hydrogen. 7) For charged molecules only (molecules in which a + or charge is shown in the formula), determine the amount of charge on each atom by subtracting the number of electrons it has from the number of valence electrons it normally has. For this step ONLY, each bond counts as one electron and each lone pair counts as two electrons. Handy hint: Hydrogen, the halogens, and carbon s family NEVER have charge, so you don t need to do this calculation for these atoms. Examples: Do ammonia, carbon dioxide, NH 4 +1 (For NH 4 +, show the math for step 7 to indicate why nitrogen has +1 charge)

11 Notes: Polarity in Covalent Compounds Quick recap: Ionic compounds form when atoms with very different electronegativities combine with each other (~1.7 or so, though this varies). This causes electrons to transfer between the atoms to fulfill the octet rule. Covalent bonds form when atoms with similar electronegativities bond with each other. Because both elements want to grab electrons, they re stuck sharing electrons to get their full octets. What happens, however, when we have two elements bond that are close in electronegativity but one still wants to grab electrons more than the other? Example: H-O bonds. H has an electronegativity of 2.2 O has an electronegativity of 3.4. The difference: 1.2 not enough to make it an ionic compound, but O is definitely different than H! In this case, the H-O bond is said to be polar covalent: Polar covalent bonds: Bonds in which the electrons aren t completely transferred from one atom to another, but where they aren t shared equally. Polar means that the electrons are unequally shared. How we show this: This sign denotes partial charge δ δ + O H The arrow points to the more electronegative atom because the electrons spend more time there.

12 In the same way that you have a polar covalent bond, you can also have a polar covalent molecule (which is usually just referred to as a polar molecule. o Polar molecules are molecules in which the electrons are unequally distributed. This is caused by the asymmetric arrangement of polar covalent bonds. Show them a model of NH 3 that describes what this looks like, and explain how the Lewis structure of NH 3 verifies this. How to tell if a molecule is polar (Example: CH 2 F 2 ) o If possible, switch the substituents on the central atom so that they are as asymmetrical as possible. ONLY do this by switching substituents, not just by moving them. o If any of the things bonded to the central atom is different than anything else, the molecule is polar. o The greater the differences in things that are stuck to the atom, the more polar the molecule. o A dipole arrow is used to point toward the most electronegative side of the molecule. Examples to do together: CH 4, NH 3, H 2 O, HF to show increasing polarity. o Show the process for drawing each of them and determining the overall polarity of the molecule. More practice examples: PF 3, NO 3-1, BF 3

13 Why polarity is important: Polar molecules have higher melting and boiling points than nonpolar molecules. o Because polar molecules have unequally distributed charge, they act like little magnets and stick together a little bit. Polar molecules dissolve well in water (and other polar liquids). o Because water is polar, polar molecules dissolve well in it because it s like mixing a bunch of little magnets together. o This is handy because many chemical reactions occur in water.