Unit 13 Review: Types of Bonding and Phase Changes (Yes, I know that the outline levels are doing stupid things partway through this document.

Transcription

1 Unit 13 Review: Types of Bonding and Phase Changes (Yes, I know that the outline levels are doing stupid things partway through this document. I m sorry about that, but I m not going to spend any more time fighting with them. Sorry.) I. Bonding A. Atoms form bonds in order to become more stable. Usually a bond forms in order to give the atoms involved a full octet of valence electrons. This more stable electron arrangement has lower energy. B. The type of bond two atoms will make is based on their electronegativity (attraction for electrons shared in the bond). 1. If the electronegativities are equal or very close to equal, the two atoms share the electrons equally. a. The bond is nonpolar or purely covalent. b. I could say the bond has 100% covalent character. (It is all covalent.) c. I could also say the bond has 0% ionic character. (It is not ionic at all.) d. Purely covalent bonds mostly occur only when the two atoms involved are both of the same element. (A oxygen-to-oxygen bond, for example, is purely covalent, AKA nonpolar covalent. ) 2. The more the electronegativities differ, the less equally the two atoms share their electrons. The more electronegative atom is an electron-hog. Because it has more ability to attract them, it keeps the electrons near it more than its fair share. a. This makes the bond polar or polar covalent. (1) The bond is polar towards the more electronegative atom. (2) The more electronegative atom ends up with a partial negative charge. This is a fancy way of saying it is a little bit negative. (3) (The other atom ends up with a partial positive charge. It is a little bit positive.) b. The bond has some covalent character and some ionic character polar covalent is the middle ground. The atoms aren t sharing equally (that would be purely covalent) and they aren t actually transferring the electrons to the more electronegative element (that would be ionic). They re in between. c. Most covalent bonds involving two different elements are this type polar covalent. 3. When the electronegativities differ by a large enough amount, the atoms give up all pretense of sharing electrons. The atom with the high electronegativity beats up the other one and steals its electrons. a. Now the electrons have been transferred. (1) The atom that stole them has extra electrons now and therefore has a negative charge. (It is an anion.) (2) The atom that lost the electrons now has a positive charge. (It is a cation.) (3) This bond has a high % ionic character and a low % covalent character. 4. If a table of electronegativities is available, I can use it to determine whether a bond is covalent, polar covalent, or ionic by finding the difference between the electronegativities of the two elements involved. a. Covalent bonds result from an electronegativity difference less than 0.5. b. Polar covalent bonds result from an electronegativity difference between 0.5 and 1.7. c. Ionic bonds result from an electronegativity difference greater than Most of the time we don t use a table of electronegativities. We make generalizations because it is quicker and easier. a. Electronegativity increases up and to the right on the periodic table the elements that are up and to the right are the nonmetals. Hence these generalizations: (1) Nonmetals have high electronegativity. (They are not willing to release their valence electrons.) (2) Metals have low electronegativity. (They release their valence electrons relatively easily.) b. Instead of looking up the electronegativities and doing the math, we usually say the following: (1) If they are both nonmetals, their electronegativities are both high (and therefore relatively similar. They will bond covalently. (If both atoms are the same element, they will bond purely covalently otherwise polar covalently.) (2) If one is a metal and one a nonmetal, the metal has a much lower electronegativity than the nonmetal. The nonmetal will probably steal electrons from the metal and the bond will be ionic. C. Covalent bonding

2 1. Is the result of electrons being shared. 2. Neither atom releases its valence electrons. The two atoms share some of their electrons in order to complete their octets. a. They are bonded to make a molecule. b. For a molecule, the molecular formula indicates the actual number of atoms that are bonded to one another. (1) Since the molecular formula is based on the actual number of atoms bonded to one another, it is not always the smallest possible ratio of the elements involved. (2) For example, glucose is covalent and has the molecular formula C 6 H 12 O 6. (a) That isn t the smallest whole number ratio for those atoms because in a molecule of glucose there are actually six carbons, twelve hydrogens and six oxygens. (b) If you simplify it (make it the smallest whole number ratio: C 1 H 2 O 1 ) it isn t the molecular formula of glucose anymore it s the empirical formula of glucose (c) (Empirical formula is a fancy name for smallest whole number ratio c. One molecule is not bonded to the next molecule. They are only attracted by IMFs which are super weak. This makes it relatively easy to separate one molecule from another, so melting and boiling points for these substances are low. D. Ionic bonding 1. Is the result of electrons being transferred. 2. The more electronegative atom takes electrons from the other atom. This forms ions which stick together because oppositely charged ions are attracted to one another. a. The ions bond in a crystal lattice b. Formulas for ionic compounds are expressed as a formula unit (1) The ionic formula is the formula of one formula unit (2) This should always be given as the smallest whole number ratio (3) That ratio is NOT the actual number of atoms that are bonded to one another. (Important difference from molecular formulas.) (4) They are making a big crystal that has the overall ratio of the formula unit but since everything in the whole crystal is bonded together, the total number of bonds actually depends on how big the crystal is... which is why we don t use that for the formula (a) It would be silly because then every different crystal of salt would have its own formula based on its size (b) Even though they were really all basically the same stuff. (Unlike covalent compounds, ionic compounds have the same properties whether they re made of ten sodiums with ten chlorines or a million sodiums with a million chlorines.) c. So in crystal structure, all the ions in the crystal are bonded together (instead of being bonded into individual molecules that don t care about each other, like the covalent compounds). That is why very high temperatures are necessary to melt ionic compounds all the bonds throughout the crystal must be broken in order to free a formula unit d. The checkerboard pattern of crystal lattice structure locks ions into place. Crystal lattice structures are very rigid because moving an ion over one row or column would mess up all the attractive forces. (1) This makes ionic compounds are very brittle (if you hit it with a hammer it will probably shatter, because moving ions messes up the attractions) (2) This makes ionic bonds non-conductors when they are in the solid state (because the ions can t move, and to conduct electricity mobile charged particles are needed) E. Metallic bonding 1. Is the result of electrons being released and delocalized in an Electron Sea. 2. All metals have relatively low electronegativities that is why they have metallic properties. Because they are NOT electron-hogs. a. The metal atoms release their valence electrons in an ionic compound, a nonmetal would TAKE those valence electrons, but when only metals are present (because it s a metallic compound) no one takes the valence electrons and they just flow around everywhere, unclaimed by any atom. (1) The electrons are considered delocalized because they are not attached to a specific location (or atom) in the metallic crystal. (2) This means the electrons can move freely throughout the entire metal sample sloshing

3 around like water in the ocean hence the electron sea moniker. b. This community pool of all the valence electrons is the cause of all the properties commonly associated with metals: (1) Luster: the electron sea gives metallic compounds a lustrous (shiny) appearance because there are electrons everywhere ready to reflect light. (2) Electrical conductivity: the electron sea makes them excellent conductors of electricity because they have so much mobile electrical charge. (3) Thermal conductivity: the electron sea can also transfer heat from one part of a metal to another. (4) Malleability and Ductility: whatever shape a metal sample is hammered (malleability) or stretched/drawn into (ductility), the electron sea will flow easily into the new shape. c. Metallic compounds bond as metallic crystals, with metal cations (+ charge because they ve lost valence electrons) in a grid, electron sea flowing around them. d. Formulas for metallic compounds are expressed as formula units and are given in lowest whole number ratio, just like metallic compounds. F. Hey, remember all those details we listed in our chart of Bonding-Type notes? You still need to know that stuff. a. Strongest bond to weakest bond: nonpolar covalent > polar covalent > ionic > metallic (all these are stronger than IMFs, which are NOT bonds) b. Lowest melting point: covalent substance; highest melting point: ionic substance c. Both ionic and metallic compounds have higher melting temperatures than covalent (molecular) compounds, because you must BREAK either ionic bonds or metallic bonds in order to melt them. (You do NOT have the break covalent bonds to melt a molecular substance you only need to overcome IMFs. The molecules separate from each other but do not break down within themselves.) II. Phase diagrams & heating curves A. Phase diagram (Pressure on y-axis, temperature on x-axis) B. States of matter 1. Solids (A on this phase diagram) a. Exist at the lowest temperatures and highest pressures* b. Have definite shape and volume c. Have low kinetic energy (particles vibrate in place but can t slide past each other) d. Have high density 2. Liquids (B on this phase diagram)

4 a. Have definite volume, but take the shape of their container b. Have medium kinetic energy (particles vibrate and can flow) c. Have medium density* d. *except water, which has a denser liquid phase, which is why ice floats! 3. Gases (C on this phase diagram) a. Exist at the highest temperatures and lowest pressures b. Take both the shape and the volume of their container c. Have high kinetic energy (particles move rapidly throughout the container; this makes gases diffuse) d. Have very low density there is a substantial distance between particles in a gas C. Phases change vocab: 1. s l = fusion or melting (latent heat of fusion is absorbed) 2. l s = crystallization or freezing (latent heat of fusion is released) 3. l g = vaporization or boiling (latent heat of vaporization is absorbed) 4. g l = condensation (latent heat of vaporization is released) 5. s g = sublimation (latent heat of sublimation is absorbed) 6. g s = deposition (latent heat of sublimation is released) 7. Triple point (E on this phase diagram) the pressure/temperature combination (on a phase diagram) where all three phases are at equilibrium all three states of matter can coexist. 8. Critical point (D on this phase diagram) the point coinciding with the critical temperature above this temperature only the gas phase can exist. (The particles have too much energy (are moving too rapidly) to stay together in the liquid or solid state.) D. Heating Curve 1. Kinetic energy (temperature) vs. Potential energy (latent heat energy of phase change) a. During the three slanted portions of the heating curve (a, c, e) kinetic energy is increasing and potential energy is constant b. During the two flat-line portions of the heating curve (b, d) kinetic energy (temperature) is constant. The energy added goes to increase the potential energy (a phase change occurs) E. Vapor pressure graph

5 1. Boiling occurs at whatever temperature causes thee substance s vapor pressure curve to cross the atmospheric pressure line. 2. At sea level, atmospheric pressure is kpa orr 1.0 atm. III. Don t forget: A. Covalent BONDS are polar when the two atoms involved in the bond don t have the same electronegativity. (A carbon to chlorine bond, for example, is polar.) B. Covalent MOLECULES are polar when polar bonds (or unshared pairs) around the central atom are not distributed symmetrically in 3-D. 1. CH 4 has polar bonds, but is a nonpolar molecule, because it is symmetrical 2. CH 3 F has polar bonds and is a POLAR molecule, because it is not symmetrical: 3. H 2 O has polar bonds and is a POLAR molecule because it is not symmetrical (remember, this is bent (109) because of the unsharedd pairs on the oxygen). 4. O 3 has nonpolar bonds (O-O) on the center atom that make it bent (120). but is a POLAR moleculee anyway because of the unshared pairs a. Also I resonated the O 3, because reasons..