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chem exam 4
| Question | Answer |
|---|---|
| who discovered valence bond theory | linus pauling and others in the 1930s |
| valence bond theory | in COVALENT BONDS, covalent bonds form when unpaired electrons, which reside in orbitals, are shared by the overlapping of single occupied orbitals bonding is based off of unpaired electrons |
| electrons in overlapping orbitals | must have opposite spins because no two electrons can have same quantum numbers (pauli exclusion?) |
| what does valence bond theory do well in explaining | the bonding in diatomic molecules with only single bonds the lack of bonding experienced by noble gases (they dont have any unpaired electrons,,,so they dont need to, and wont, bond) |
| implications of valence bond theory | to form a covalent bond, an atom must have an unpaired electron (if electrons are all paired off, they have no incentive) the number of bonds formed by an atom is determined by the number of unpaired electrons |
| issue with valence bond theory | it must be modified to explain the covalent bonds found in other molecules (like BF3 and CH4) |
| what energy states are more stable | lower energy states |
| why do atoms bond? in terms of energy | to get to lower, and therefore more stable,energy states |
| qhat determines bond length? | the internuclear distance at which the energy is lowest for the two "optimum distance to achieve lowest overall energy of system" |
| hybridization | the mixing of two or more atomic orbitals to form a new set of hybrid orbitals mix at least 2 nonequivalent atomic orbitals (such as s and p) |
| orbital | a region in space an electron is likely to be found (one of the circles in an orbital diagram is one orbital) |
| subshell | grouping of orbitals (the circles in the diagram grouped; 3 circles for p, 5 for d, 1 for s, 7 for f) that share the same angular momentum quantum number (l) |
| hybrid orbital shapes | very different from their respective original atomic orbitals ex: sp different than s and p shapes; new third shape |
| how can you tell how many hybrid orbitals you have | the number of hybrid orbitals is equal to the number of pure atomic orbitals used in the hybridization process ex: if you used an s subshell and a p subshell, thats 4 orbitals total (1 from s and 3 from p), so youll have 4 hybrid orbitals |
| two ways covalent bonds are formed | 1) overlap of hybrid orbitals with ATOMIC ORBITALS 2) overlap of hybrid orbitals with OTHER HYBRID ORBITALS |
| what orbitals are being hybridized/combined in hybridization | orbitals in the same atom, so they can make more unpaired electrons and make more bonds ex: carbon has both s and p orbitals, it will hybridize them in a way so that it has a total of 4 unpaired electrons, which is why CH4 can be formed |
| goal of hybridization | to get more unpaired electrons so you can form more bonds |
| sigma bonds | DIRECT overlap of ANY two orbitals |
| pi bonds | SIDE BY SIDE overlap of two P ORBITALS think pi p |
| what types of bonds does hybridization happen with | covalent not ionic for example |
| bonding in other molecules best explained by | hybrid orbitals |
| conservation in hybridization | the # of hybrid orbitals formed always equals the # of atomic orbitals that are combined |
| need hybrid orbitals to explain | geometry |
| one region of electron density | could be a lone pair, unpaired electron, single bond, or multiple bond all just 1 region of electron density |
| single bonds are formed by | the direct overlap of 2 hybrid orbitals, p orbitals, or s orbitals sigma bonds |
| additional electrons shared in a multi bond formed by/result of | not a result of directly overlapping hybrid orbitals like single, but rather a result of side-by-side overlap of 2 regular "p" atomic orbitals (unused in hybridization for example) |
| what kinds of bonds are in a triple bond | 1 sigma and 2 p bonds |
| in a carbon atom, for example, if you only use 2 of the p orbitals for sp2, how do you decide which p is not used | its pz, because it needs to be able to stick out of plane to form the pi bonds with another carbon atom px and py are within the plane pz sticks out above and below the plane |
| what plane are we looking at when looking at orientation of orbitals and sticking out and whatnot | sigma system plane all sigma bonds will be in that plane; all ppi bonds will be perpendicular to it |
| what type of bond is formed when a hybrid orbital overlaps a hybrid orbital | sigma |
| are sigma or pi bonds higher in energy | pi think; double and triple bonds have higher energy than single and they have pi bonds |
| how to determine how many hybridized orbitals you need | look at how many areas of ELECTRON DENSITY (not just bonds) are around the atom if 2, youll do sp and the other 2 ps will be normal and if connected to another atom like this, those 4 ps will form 2 pi bonds total |
| do you always have to use the max amount of hybrid bonds that CAN be formed | no for carbon for example max we can make is sp3,,,doesnt mean we have to we could also have sp2 and 1 p normal or sp and 2 normal p orbitals |
| what does molecular orbital theory try to explain | tries to explain the difference between what we see/predict in theory and what we see in experiment for example, O2 should have no unpaired electrons, so it should be dimagnetic, but in experiments sometimes it has reaction to magnetic field |
| molecular orbital theory | bonds are formed from interaction of atomic orbitals to form molecular orbitals |
| bonding molecular orbital | has lower energy and greater stability than the atomic orbitals from which it was formed where we see good bonds between molecular orbitals |
| antibonding molecular orbital | has higher energy and lower stability than the atomic orbitals from which it was formed represented by star or asterisk in bonding diagram |
| if you have higher energy, you have ____ stability energy and stability are related how? | lower inversely think:because everything in world wants lower energy,,,you have low energy youre content youre stable you have no reason to react |
| construction interaction of 2 hydrogen molecules | bonding sigma molecular orbital means they bond |
| destructive interaction of 2 H molecules | antibonding sigma molecular orbital they...repel? can bond in an antibonding orbital |
| bond order | 1/2 (number of ELECTRONS (not electron PAIRS) in bonding molecular orbitals - number of ELECTRONS in antibonding molecular orbitals) |
| why do we see constructive and destructive behavior in atoms bonding | because of the wave nature of electrons atoms electrons are behaving this way, and waves can either be destructive (cancel out) or constructive (add together) |
| molecular orbital v atomic orbital | |
| boyles law | the volume of a sample of gas at constant temperature varies INVERSELY with the applied pressure PV=PV |
| charles's law | the volume of a sample of gas at constant pressure is DIRECTLY PROPORTIONAL to the absolute temperature (IN K) |
| absolute zero | -273.15 K the temp @ which volume of a gas is hypothetically zero can only have a volume of zero if its not moving at all (temp is a measure of kinetic energy in this case, no temperature means its not moving &if its not moving its not taking up space_ |
| kinetic-molecular theory (kinetic theory) | a theory, developed by physicists, that is based on the assumption that a gas consists of molecules in constant random motion |
| molecular speeds | according to kinetic theory, molecular speeds vary over a wide range of values. The distribution depends on temperature, so MOLECULAR SPEED INCREASES AND TEMPERATURE INCREASES |
| root-mean square (rms) molecular speed, u | a type of average molecular speed, equal to the speed of a molecule that has the average molecular kinetic energy [insert eqn] |
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