CHAPTER 6 CHEMICAL BONDING. Chemical Bond a link between atoms that holds them together in a compound.

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1 CHAPTER 6 CHEMICAL BONDING Chemical Bond a link between atoms that holds them together in a compound. Why Bonding Occurs usually to get to a lower energy state. Most atoms drop in energy when they form bonds with other atoms, because bonding allows them to get to an octet. **OCTET RULE most atoms will gain lose, or share enough e - to form an octet (8 e - ) in the outer shell. Exceptions: 1. H and He want 2 e - since their outer shell only holds 2 e-. 2. Some atoms form less than or more than an octet due to unusual bonding patterns. Don t need to know details. **Since forming a bond causes a drop in energy, most bond formation is exothermic (the atoms release E as they bond). Similarly, bond breaking is usually endothermic. I. TYPES OF BONDS A. Ionic bond one atom loses e - to get to an octet (producing a cation), another atom gains e - to get to an octet (producing an anion). The ions are then attracted to each other by their opposite charges. EX: A + B A + + B - AB B. Covalent Bond results from 2 atoms sharing electrons, so both atoms can have an octet. Ex: A + B AB 1

2 Bonds are rarely completely ionic or completely covalent. They are somewhere in a range. The type of bond formed between any 2 atoms can be determined by subtracting their electronegativities. Non- Ionic Bond Polar Covalent Polar completely completely ionic covalent Nonpolar covalent bond the bonding e - s are shared equally between the two atoms; there is a balanced distribution of charge. A B Polar covalent bond e - are shared, but somewhat unequally. One atom pulls the e - harder than the other does. 2

3 PRACTICE Cs F bond: = 3.3 Ionic M-N C O bond: = 1.0 Polar Covalent N-N P Cl bond: Polar Covalent N-N H Br bond: Polar Covalent N-N Li O bond: Ionic M-N S Se bond: Non Polar Covalent N-N Trends: metal + nonmetal = Ionic nonmetal + nonmetal = Covalent Quiz #1 Types of bonds 3

4 II. Metallic Bonding occurs in solid or liquid metals (Cu, Hg, Ag, Au, Ni, Na, etc.) These atoms give up e - to get to an octet, but the e - don t go to another atom. Instead, the e - move freely around the sample of metal. A. Electron-Sea Model the piece of metal is held together by the attraction between the cations and the sea of electrons. Ex: B. Properties of Metals all explained by the electron-sea model of metallic bonding. 1. Good conductors of heat and elec. because the e - are free to move, transporting the charge (or the heat) to the other end of the substance. (True in all phases). 4

5 2. Malleable and ductile because the e - sea operates within a plane or layer. So the bonding is: Metal Ionic Cmpd 5

6 3. Shiny because the e - sea allows each e - to have a wide range of possible energies (they aren t locked into the energy levels of a particular atom), so they can absorb and re-emit light of many wavelengths. 4. Bond Strength directly proportional to the # of e - each metal atom gives up. (A cation with higher charge, such as +2 or +3, pulls harder on the sea of e- than a cation with charge of +1). Ex: Na +1 vs. Fe +2 or Fe +3 (soft, cut with a knife) (very hard metal) 6

7 III. IONIC BONDING AND IONIC COMPOUNDS Ionic bonds: metal + nonmetal A. Formation: 1. ions form 2. they attract each other due to charge **Determining what ion an element will form: 1. Determine the # of valence e - 2. Decide how many it needs to gain or lose to get to eight (octet rule). It always chooses the path (gaining or losing) that involves the fewest electrons. 3. Since the + and charges are no longer balanced, the neutral atom has become an ion (charged). If it gained e -, the final charge will be negative. If it lost e-, the charge will be positive. Practice with disks, using Na, Cl as examples. Na: has _11 electrons. Has _1 valence e -. To get to octet, could gain 7 e -, or lose 1 e -. Will _lose 1 e -, giving a charge of _ + 1. The ion is _Na +1. Cl: has 7 valence e -. To get an octet, can gain _1 e - or lose 7 e -. After gaining 1 e -, the ion formed is: _Cl -1. Al: Will lose 3e - to get to octet. The ion formed is Al +3. O: Will gain 2e - to get to octet. The ion formed is O -2. N: Will gain 3e - to get to octet. The ion formed is N -3. Ca: Will lose 2e - to get to octet. The ion formed is Ca +2. S: Will gain 2e - to get to octet. The ion formed is S -2. 7

8 A. Ionic Cmpds composed of positive and negative ions drawn together by their charges in a ratio that balances the + and charges to give a neutral compound. **Most ionic cmpds are 3-D crystals. --this is the lowest energy position --So, the formula for an ionic compound does not represent one molecule (an independent particle), but represents the simplest ratio of ions that gives electrical neutrality. Ex: NaCl doesn t mean one molecule of NaCl, but shows that a ratio of one Na + and one Cl - gives neutrality. Ex: What is the most likely formula for a compound of Al and Br? Al forms Al +3 ion Br forms Br -1 ion To balance the charges takes 1 Al +3 and 3 Br -1 ions. Formula: AlBr 3 (**always write the cation first) Practice: What is the most likely formula for a compound of Mg and S? MgS A compound of I and Na? NaI A compound of Ca and N? Ca 3 N 2 A compound of Pb +4 and O -2? PbO 2 Quiz #2 Ionic and Metallic compounds 8

9 B. Comparing Ionic and Covalent Compounds 1. Comparing Bond Strength a. ionic bonds between ions within the crystal. Are fairly strong due to the attraction between oppositely charged ions. b. covalent bonds between atoms within a mcl. Are very strong, because breaking them apart causes each atom to lose its octet. c. intermolecular forces (IMF s) forces of attraction from one covalent molecule to another. Are fairly weak. (ex: water molecules hold together into a drop, but not strongly). 2. Comparing Properties a. Solubility Ionic cmpds are generally more soluble in water than covalent cmpds. b. Melting Point/Boiling Point For a solid to melt, the particles have to be separated so they can move freely in the container. Is that easier for an ionic cmpd, or a covalent one? --For an ionic cmpd, you must break the bonds between ions in an ionic crystal; these forces are strong. -- For a covalent cmpd, you must break the IMF s between adjacent covalent mcls. These forces are _weak. So covalent cmpds generally have lower MP and BP than ionic cmpds. (melt and boil more easily) c. Hardness 1. Ionic ions are held strongly in place by ionic bonds. Very hard. 2. Covalent mcls held in place by weak IMF s. Not as hard as ionic cmpds. 9

10 d. Conductivity how easily electric current passes through a substance. Electric Current caused by the flow of charged particles through a substance. So, for current to be conducted, you must: (1) have charged particles, and (2) they must be able to flow. 1. Ionic cmpds Solid: ions, not free to move No conductivity Liquid: ions, free to move Conductivity Dissolved in H 2 O: ions, free to move Conductivity 2. Covalent Cmpds Solid: no ions Liquid: no ions Dissolved in H 2 O: no ions IV. COVALENT BONDING AND MOLECULAR COMPOUNDS ( Molecular cmpd is the same as covalent cmpd ) A. Molecule (mcl) the smallest particle of a cmpd that can exist while still having the properties of that cmpd. Chemical Formula tells the relative numbers of each kind of atom found in a certain compound. Ex: H 2 O: 2H atoms, 1 O atom per mcl C 2 H 5 OH: 2 C, 6 H, 1 O per molecule Ca(OH) 2 : 1 Ca, 2 O, 2 H ratio in the crystal 10

11 B. Formation of Covalent Bonds Ex: H-H bond Each H atom: H H As they come close, 1. the e- repel each other 2. each nucleus attracts the other atom s e- (stronger force) So, they are drawn together. The e- clouds merge. H H If they get too close, the nuclei start to repel each other. Finally the attractive and repulsive forces balance each other, and the atoms settle into being a molecule. 1. Bond length the average distance between the nuclei of two bonded atoms. Bond lengths generally increase when larger atoms are involved. 2. Bond energy the E required to break a chemical bond and form neutral atoms. **CLASS ACTIVITY : Gumdrop molecules Needed per group: 5 toothpicks Ruler 3 small gumdrops (same color) 3 large gumdrops (diff. colors) Molecule #1: 1 toothpick with 1 small gumdrop at one end, 1 large gumdrop at other end. Molecule #2: 2 toothpicks with 2 cm broken off one end of each; hold them together, put 1 small and 1 large gumdrop on ends. Molecule #3: 2 toothpicks broken in half; hold all four pieces in a bundle together, put 1 small and 1 large gumdrop on ends. 11

12 You have built HBr, HI, and HCl. Based on their positions on the P.T. and the new info about bond length, which mcl is which? Why? Try to break the bonds joining the atoms. Trend: As bond length increases, bond energy. (atoms are further apart, so they can t hold as tightly) C. Drawing Covalent Molecular Structures (Lewis Structures) --use the octet rule and e- dot notation. --e- dots can be shifted around the atom if needed to help arrange the atoms. Ex: Structure of H 2 mcl H H H H Ex: Cl 2 mcl Cl Cl Cl Cl unshared pairs shared pair Unshared pair two e- that are not involved in the bond; instead they remain exclusively with one atom. In a Lewis Structure, the shared pair can be replaced by a dash. H-H or Cl-Cl (this is a timesaver in complex mcls) 12

13 Steps in Drawing Lewis Structures 1. Determine the type and # of atoms in the mcl (may be given, or may need to find based on # of bonds each wants to form). 2. Draw the e- dot symbols separately for each atom. **Each unpaired e- dot should be thought of as an attachment point for another atom. 3. Arrange the atoms next to each other to produce an octet in each one (except H and He). *Note: In a multi-atom mcl, the atom that is closest to having four valence e- is placed in the center, because it needs the most bonds to get an octet. Ex: Draw the Lewis structure for CH 3 I Step 1: Step 2: put e- dots around them. Step 3: C is closest to having 4 valence e-, so put it in the center. Ex: Draw the Lewis Structure for NH 3 13

14 Ex: Draw athe L.S. of a cmpd of Si and F Step 1: Si has valence e- Si (4 attschment pts) F has valence e- F (1 attachment pt) So, need: Independent Practice: Draw the L.S. s for: 1. PCl 3 2. A cmpd of H and S 14

15 Single bond a covalent bond in which one pair of e- is being shared. (all covalent bonds shown so far are single bonds) D. Multiple Covalent Bonds --some elements, especially C, N, and O, can share more than one pair of e- with another atom. Double bond when two pairs of e- are shared between 2 atoms. Ex: Ethene Triple bond when three pairs of e- are shared between 2 atoms. Ex: N 2 Drawing Lewis Structures w/multiple Bonds --Remember that C, N, and O will often form multiple bonds if needed to get an octet. Ex: Draw the L.S. for CH 2 O (formaldehyde) So, the unpaired e- join to form a double bond. Independent Practice: Draw L.S. s for: 1. CO 2 HCN Quiz #3 Drawing Lewis Structures 15

16 E. Diatomic Elements always found in nature as a double-atom mcl; covalently bonded. Never found as a single, free atom. There are seven diatomic elements know these. H 2 N 2 O 2 F 2 Cl 2 Br 2 I 2 H-H N N O=O F-F Cl-Cl Br-Br I-I V. MOLECULAR GEOMETRY --the properties of mcls depend not only on bonding (ionic/covalent/metallic), but also on the shape of the mcls. A. VSEPR Theory used to predict the shapes of covalent mcls. Valence Shell Electron Pair Repulsion Theory Electron pairs (shared or unshared) will repel each other (same charge). So, they orient themselves in space so they will be as far apart as possible. The five most common shapes: 1. Linear ex: H 2 (H-H) HCl (H-Cl) CO 2 (O=C=O) 2. Tetrahedral ex: CH 4 Lewis Structure: To get the e- pairs as far apart as possible, they form a tetrahedral shape. (Model) 16

17 3. Triangular Pyramidal like tetrahedral, with an unshared pair at the 4 th point of the tetrahedron. Ex: NH 3 Lewis Structure: Shape: 4. Bent like a tetrahedral, but with two unshared pairs at two points of the tetrahedron. Ex: H 2 O Lewis Structure: Shape: 17

18 5. Triangular Planar ex: H 2 CO (formaldehyde) Lewis Structure: Model? To get these as far apart as possible, they form a triangle. B. Steps for Predicting What Shape a Molecule Will be: 1. Draw the Lewis Structure 2. Count the # of atoms and the # of unshared pairs coming off the central atom. 3. If the total is : 1 or 2: the molecule is linear. 3: the basic shape is triangular planar.* 4: the basic shape is tetrahedral.* *If the basic shape is triangular planar or tetrahedral, check if any of the items counted in Step #2 were unshared pairs. If so mentally cover those corners of the tetrahedral (or triangle) and see what shape remains (triangular pyramidal or bent). Practice: Determine the shaped of the following molecules: 1. SF 2 2. AsI 3 3. HCN 18

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