Noble Gases are the most elements. Why? Notice that this makes a full outer energy level have electrons.

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1 NAME: Mods: Now that we know proper formula writing and naming of chemical compounds so we can speak the language of Chemistry, let s move on to understanding how and why these compounds are put together! So, let s get started by getting some info on IONIC BONDING! 1. Remember that the Noble Gases are the least reactive elements (also called the Inert Gases), and that the other elements are on a Quest for Nobility. This is because the Noble Gases are the most elements. Why? Besides Helium, the Noble Gases have this electron configuration: Notice that this makes a full outer energy level have electrons. These outer energy level electrons are known as electrons. 2. IONIC BONDING occurring mainly between a and a. Our most reactive metals are the Alkali Metals, and nonmetals, the Halogens. Let s see why they like to bond with each other: How many valence electrons do the alkali metals have? How many valence electrons do the halogens have? Look at the drawing of the atoms below: What do you notice about the sodium atom s 2 nd energy level? How many electrons away is the chlorine atom from having a full outer energy level?

2 3. Both sodium and chlorine are very reactive, and when put in contact with one another, there is a very violent reaction releasing a lot of heat and light. However, what is left behind is a white solid that we know as sodium chloride (NaCl). Why is salt not as dangerous to have around the house as the elements that make it up? 4. So, how was this sodium chloride made? Sodium readily gave up its single valence electron and chlorine readily accepts it. When this happens, what happens to the atoms of these elements? Na becomes Cl becomes The force of attraction between the +1 charge of the sodium cation and the -1 charge of the chloride anion creates a what? While we commonly call sodium chloride, chemists actually use that term to mean what? This transfer of electrons happens on a very large scale so that each cation and anion attract several of the opposite ions, creating the very distinctive structure that all ionic compounds have. Also, this transfer of electrons involves energy changes. Here is how it looks from an energy perspective:

3 5. Recall that our ionic compounds (salts) are made of and, and they are held together by. So, what else? Ionic compounds DO NOT consist of! (!) Both and forces exist within a salt crystal, but the attractive forces are significantly greater. When a salt melts (is MOLTEN) or dissolves in water, what happens? Most ionic compounds are both and. 6. There is a special type of bonding that exists among metal atoms that accounts for several of the properties of metals. Of course we already know that metals will always lose electrons to form cations, and we should know that they are doing this to become more stable (to achieve stable electron configurations). But what if there are no nonmetals available to accept these electrons? o This pool of electrons allows metals to and accounts for its. The metallic bonds in some metals are stronger than others, based on how many are being shared. For example, the alkali metals only have one valence electron, making their bonds relatively : they are soft enough to cut with a knife and melt at relatively low temperatures (sodium at ). Transition metals, like tungsten, have more valence electrons and are harder and melt at higher temperatures (tungsten at ). o Metals will also form alloys:

4 o Alloys have revolutionized our world, starting with the Bronze Age all the way up to our high-tech planes, missiles and spacecraft. Here are some common alloys: 1. GOLD: Pure gold is -karat, but very soft and can be easily wornaway or dented. So, gold can be alloyed to make it stronger. 12-karat gold is gold while 18-karat gold is gold. 2. BRONZE: Made with and. These two metals are relatively soft, but much harder and stronger when mixed. 3. BRASS: Made with and. Brass is softer than bronze, easier to shape and shinier (used to make ). 4. STEEL: An alloy of containing from %. The carbon atoms form bonds with the iron atoms making it stronger. Stainless steel contains almost no carbon, but is 10%. Steel cables on bridges (like the Golden Gate Bridge) must be able to stretch, so they also contain,, and. 5. SUPERALLOYS: Have revolutionized the aircraft industry. Small amounts of is added to aluminum and magnesium allowing them to maintain their strength at very high temperatures (like in ). 7. Now it s time for COVALENT BONDING! In ionic bonding, there was a transfer of electrons from one atom to another. With covalent bonding, our neutral nonmetal atoms are going to electrons! Sharing is caring! Yeah Covalent Bonding!

5 The simplest example of sharing electrons is found in our diatomic molecules, and the simplest of those is the Hydrogen molecule! Let s look at how the bond is formed between two Hydrogen atoms: Why is this happening? What does each Hydrogen atom now resemble? What is this sharing of electrons known as, and what do we call the space that these electrons occupy? How do we know that the molecule is more stable? Remember potential energy? Notice that the bond length (the distance between the nuclei of the atoms) is (picometers). This is an average distance, as bonds are more like a spring than a rigid pole and they vibrate, varying in length. The shorter the bond, the it is, and the more energy it will take to break it (BOND ENERGY). Here are some examples:

6 8. Remember electronegativity? We said it was how bad an atom wants an electron and that FLUORINE was our most electronegative element. This makes sense Fluorine, along with the other halogens, really want one more electron to complete their octet. In general are much more electronegative than. Here are the electronegativites of several elements: So, why are these important? Well, just like with humans, not all sharing of electrons between atoms is done equally some atoms want them more! So, these types of covalent bonds arise: NONPOLAR COVALENT BOND: POLAR COVALENT BOND: Dipole: So how do we know what type of bond will form between elements?

7 Notice that the larger the difference between the electronegativities, the the bond. Also notice the polar covalent electron cloud picture. The dipole is represented by the symbols δ + and δ - meaning positive or negative. *(δ is the lowercase Greek letter delta) So, what type of bond will form between these elements? Ionic, Nonpolar Covalent or Polar Covalent Be and F N and O Li and H Cl and Cl Ge and S K and F Drawing Bonds 9. G.N. Lewis gave us a nice way of modeling covalent bonds that is based upon our knowledge of valence electrons: Lewis Structures. An element s valence electrons are represented by dots around the symbol of the atom. We ll start by taking a look at the Lewis Dot Diagrams for several of the Representative Elements. Yeah Bond Drawing!

8 Elements with an octet of valence electrons (like the Noble Gases) are stable, and remember that our elements share electrons to become stable! We can see through Lewis structures why Chlorine is always found as a diatomic molecule in nature: One chlorine atom: Unstable Two chlorine atoms: Stable! So, how do we draw these Lewis Structures for more molecules? Let s see some samples: Covalent Bonds: H 2 O NH 3 *By the way, when there are remaining electrons and you fill them into the drawing in pairs, those are called of electrons, and they have an effect on both the of molecules as well as some of their properties!

9 Single covalent bonds are NOT the only types of covalent bonds that exist. Perhaps you have heard of double or even triple bonds. Here are a few examples of them, which will be helpful during our molecule building and drawing activities. Double Bond (Oxygen Molecule) Triple Bond (Nitrogen Molecule) *Now, Practice! Complete the Covalent Bond Drawing Practice Sheet. 10. On a final note, it is important to remember that atoms and molecules are 3-dimensional (not just flat and motionless like on your paper when you draw them!). Molecules have very distinctive shapes and features that give the molecular compounds their properties. Like with many chemical properties, these come from the internal electrostatic forces between the like and opposite charges (protons, electrons, etc.). This leads to the VSEPR Theory: V S E P R This theory explains to us that the most stable arrangement of a molecule keeps the electron groups (atoms or lone pairs) as far away from each other as possible! Depending on how many atoms and lone pairs are found around a CENTRAL ATOM tells us what shape that molecule is around that central atom. These are five possible shapes, and they are summarized as follows: Around the central atom! Number of Atoms Number of Lone Pairs Shape around Central Atom Linear Trigonal Planar Tetrahedral Trigonal Pyramidal Bent

10 This VSEPR theory also helps determine whether or not a molecule will be POLAR or NONPOLAR. A molecule can have polar bonds, but still be nonpolar if the bonds cancel each other out. If they don t, then the molecule will be polar. Here are the guidelines for molecular polarity: 1. When a molecule contains zero or one polar bond, then the molecule with no polar bonds is a nonpolar molecule, while the molecule with one polar bond is a polar molecule. Example: NONPOLAR CH 4 POLAR CH 3 Cl 2. When a molecule has more than one polar bond, the shape of the molecule will determine the polarity of the molecule: If the individual bond dipoles do not cancel, the molecule is polar. If the individual bond dipoles do cancel, the molecule is nonpolar. Examples: We will now practice building models of molecules and determining their shape and polarity! I know, SCORE!