Test Review for Chem Test 3
Electronegativity; defined as the ability of an element to attract electrons towards itself in a bond. A higher electronegativity means a stronger attraction
∙ Fluorine is the most electronegative element at 4.0
∙ The trend increases up the columns and increases across the group (↗) ∙ Know specifically the following elements
o H (hydrogen) at 2.1
o B (boron) at 2.0
o C (carbon) at 2.5
o N (nitrogen) at 3.0
o O (oxygen) at 3.5
o F (fluorine) at 4.0
o S (sulfur) at 2.5
o Cl (chlorine) at 3.0
o Br (bromine) at 2.96
o I (iodine) at 2.66
∙ When electronegativity difference between two elements is large, the bond is polar.
∙ ΔEN=EN(element 1) – EN(element 2)
o That formula literally just says that the difference in electronegativity of the molecule is the electronegativity of the first element minus the electronegativity of the second element.
Note; element 1 is always the element with the highest
∙ A ‘dipole moment’ is when the positive and the negative charges separate. It forms a light bulb looking electron cloud. The bulb part of the light bulb is around the more electronegative atom.
∙ In ionic bonds, electrons are taken, not shared, like they are in covalent bonds. These bonds are between metals and nonmetals. Ionic bonds happen when ΔEN is greater than 2.0
∙ In covalent bonds, electrons are shared between two elements, not taken, like they are in ionic bonds. ΔEN must be similar. These bonds are between two nonmetals or between a metalloid and a nonmetal. The metals don’t know how to play nice and can’t share anything. Jerks. Don't forget about the age old question of What are the three models of communication?
∙ Pure covalent bonds have a completely even sharing of electrons. ΔEN=0. For example, Cl2, or any other diatonic element. Don't forget about the age old question of What are the main features of the heart?
∙ Nonpolar covalent bonds have a mostly even sharing of electrons. ΔEN is between 0 and .4. For example, H2O (ΔEN=1.4)
∙ Polar covalent bonds have a decidedly uneven sharing of electrons. ΔEN is between .4 and 2.0 Polar covalent molecules will have a dipole moment.
Lewis Dot Bond Theory
∙ Bonds form between elements to help elements achieve noble gas electron configurations. Elements are all trying to have an octet of valence electrons, or in hydrogen’s case, a duet of valence electrons.
o Rules of Connectivity
o 1) Hydrogen has ‘exterior terminal connectivity’, meaning when drawing models, you always stick hydrogen on the outside of the structure and don’t connect it to anything else.
o 2) The more electronegative an element is, the further it is from the center of the structure.
o 3) If a molecule has both hydrogen and oxygen, the two must go next to each other. We also discuss several other topics like What is nullification crisis?
Unless it’s an organic compound (anything with carbon, oxygen, and hydrogen)
o If those rules still can’t get all the elements to a full octet, move a lone pair from one element to a bonding pair between it and the element under octet.
∙ Steps to form bonds using Lewis Dot Structures
∙ 1) check that you’ve followed the rules of connectivity
∙ 2) count up the total valence electrons
∙ 3) distribute the total valence electrons so all elements are at octet ∙ 4) if elements are under octet, move a lone pair of electrons to become a bonding pair
∙ 5) replace all bonding electrons with bonds (ie lines) show lone pairs, bracket ions, and add up the charges of each element
∙ 6) Charge = #valence electrons brought – nonbonding electrons – ½ of the bonding electrons.
o The perfect Lewis Dot Structure will have formal charges near zero for each element. Any negative charges that can’t be avoided are with the most electronegative elements.
When trying to determine which Lewis Dot Structure is the best Lewis Dot Structure, go for the structure that has the lowest formal charge. Formal negative charges should be on the elements that are the most electronegative.
Exceptions to the Octet Rule
∙ “Odd Electron Species” means that the total electron valence count is an odd number. You end up with a radical (a single, lone, unpaired electron) on one of the elements. For example, NO2 has an extra valence electron. Don't forget about the age old question of What is relationship according to other theories?
o Reminder: electrons are distributed first to bonding pairs, then to lone pairs on exterior element, and then to lone pairs on the central element. Once all of the electrons have been distributed, you can move bonds around to reduce formal charges, if necessary.
∙ Some elements (I’m looking at you, Group 3A) prefer to be under the octet. (Specifically for group 3A, at 6 valence electrons) for example, BF6. In BF6, Boron likes 6 valence electrons because it ends up with a formal charge of zero. We also discuss several other topics like What is high self-efficacy?
We also discuss several other topics like What does the law of diminishing marginal productivity state?
∙ Some elements (cough cough periods 3-7 cough cough) can have expanded octets, meaning they can have more than 8 valence electrons. This is because they have room d shells. They must be surrounded by at least 2 other elements. This is to decrease the formal charges. For example, AsF5.
o Also, sometimes all the atoms get an octet, but there are still leftover valence electrons to be distributed. The extra valence electrons have to end up on atoms that can exceed the octet. For example, in XeF4, Xe gets 12 valence electrons.
∙ Bond energy is defined as the energy it takes to break 1 mole of a bond in a gas phase, and it’s measured in kilojoules per mole, kj/mol. Splitting a bond creates 2 free radicals.
o The bond energy depends on the strength of the bond (triple is stronger than double, double is stronger than single)
o The bond length also depends on the single/double/triple status of the bond. If you think about bond length as a certain amount of rope that stays constant, it’s easy to remember. Obviously, if you fold the rope in half for a double bond, it’s going to be shorter than if you left the rope out long, like a single bond. And if you fold the rope in three, for a triple bond, it’ll be even shorter.
∙ Stands for Valence Shell Electron Pair Repulsion Theory
∙ It gives you the 3D shape of the molecules
∙ It works because electrons repel other electrons. They all want to be as far away from each other as possible.
∙ The exception to electron spreading is pairs of electrons. Pairs of electrons can’t separate. Two different sets of paired electrons can separate, but the individual electrons in each pair can’t separate.
o This applies to electron bonds as well. Each bond counts as a group, and groups want to be as far away from each other as possible. ∙ The number of electron groups determines the shape.
∙ Defined as shape around an atom considering all electron groups ∙ Focuses on the central element (the element with 2+ groups around it) o 2 electron groups = linear, 180 bond angles
o 3 electron groups = trigonal planar, 120 bond angles
o 4 electron groups = tetrahedral, 109.5 bond angles
o 5 electron groups = trigonal bipyramidal
2 types of groups, equatorial and axial. Equatorial groups are essentially trigonal planar, but it has one group directly above
the central atom and one below, which are the axial groups
Axial to axial bonding angles are 180
Equatorial bonding angles are 120
o 6 electron groups = octahedral, 90 bond angles
∙ Based on the number of electron groups
o 2 electron groups = linear = ideal bond angles of 180
o 3 electron groups = trigonal planar = ideal bond angles of 120 o 4 electron groups = tetrahedral = ideal bond angles of 109.5 o 5 electron groups = trigonal bipyramidal = ideal bond angels of 90 & 120
Trigonal bipyramidal have 2 types of bonds, axial and equatorial ∙ Equatorial are all on the same plane, they’re basically just
trigonal planar (the 120 degree ones)
∙ Axial are stuck on the top like a radio antenna (the 90
o 6 electron groups = octahedral = ideal bond angles of 90
∙ The Lone Pair Effect
o Lone pairs count as an electron group
o They like to spread out around an atom. Bond pairs do not. Bond pairs stay pretty concentrated.
o Lone pair/lone pair repulsion is much stronger than lone pair/bonding pair repulsion, which is all still stronger than bonding pair/bonding pair repulsion.
o This means when a central atom has a lone pair, the bond angles shrink away from it. This decreases the ideal bond angles.
∙ Based on the number of electron groups as well as the number of bonds ∙ Example of a bent molecules; NO2-
∙ Example of a trigonal planar molecules; BF3
∙ Example of a tetrahedral molecule; CH4
∙ Example of a trigonal bipyramidal molecule; PCl5
∙ Example of an octahedral molecule; BrF5
Electron Geometry (EG) = Molecular Geometry (MG) if there are no lone pairs
EG =/= MG if there are any lone pairs
Non Polar v Polar Molecules
∙ The difference between the electronegativity values for each element determines the polarity of the molecule.
∙ Dipoles point towards the more electronegative element.
∙ When summing dipoles to determine polarity strength and direction, sum the dipoles up two at a time. ↗+↙=0 ←+↓=↙ →+↑=↗ ↓+↑=0
o Note; she has said we will only need to be able to approximate the direction, not the exact value or angle.
∙ A molecule can have all polar bonds and still be non-polar, if all the polar bonds cancel one another out.
∙ Polar and non-polar molecules will not mix
∙ Quick Hint; tetrahedral like CF4, where there’s only 1 type of bond, (in this case C-F) will almost always be non-polar because the bonds will cancel out.
In Lewis Dot Theory, valence electrons are represented by dots
In VSEPR Theory, you know the basic shape and bond angels of molecules In Valence Bonding Theory, you know which orbitals are used to form bonds Valence Bonding Theory
∙ Overlapping orbitals form bonds
∙ Each of the overlapping orbitals must have a single electron
∙ Each of these electrons must have opposite spins
∙ In order to get the correct bond angles for the molecular geometry of a molecule, you have to have hybrid orbitals.
o VSEPR Theory tells you which shapes and bond angles you need ∙ Atoms can only use their own orbitals when hybridizing.
o Normal P orbitals look like dumbbells. Hybrid P orbitals look like birthday balloons.
Steps of Hybridization
∙ The number of orbitals hybridized is the total number of orbitals, for example if you have 1 s orbital and 3 p orbitals, you have a total of 4 orbitals. ∙ The shape and identity of the orbitals depends on how many s/p/d/f orbitals were hybridized, for example 1 s orbital and 3 p orbitals is identified as 4 sp^3 orbitals. Only the orbitals in the principal shell (that’s the one the valance electrons are in) can hybridize.
o NOTE; the particular principal shell number does not effect the identity of the hybrid orbital. Meaning, 2 p orbitals from principal shell 3 are named exactly the same as 2 p orbitals from principal shell 2. It’s the number of p orbitals that matters, not where they came from.
∙ The hybridization depends on the number of electron groups AND the electron geometry of the atom. For example,
2 electron groups
Hybrid Orbital Name; 2 sp
3 electron groups
Hybrid Orbital Name; 3 sp^2
4 electron groups
Hybrid Orbital Name; 4 sp^3
5 electron groups
Hybrid Orbital Name; 5 sp^3d
∙ Notice how the number of electron groups tells you the hybrid orbital name. There’s only 1 s orbital, and only 3 p orbitals. You can’t have a p^4, or an s^2. It’s like doing electron configurations again. If you have 5 electron groups, you have to go all the way to your d orbitals because there’s only 1 s and 3 ps, that only gives you 4 electron groups. To get to 5, you have to move to the next level of orbitals.
Bonds determine if hybrid orbitals are used, or if non-hybrid orbitals are used.
∙ Sigma bonds (σ) are formed by overlapping of orbitals along a bond axis. Sigma bonds always require hybridized orbitals. Single bonds are always sigma bonds (σ). It alliterates, remember it. Sigma
bonds are also stronger than pi bonds.
∙ Pi bonds (π) are formed when orbitals overlap side by side. They require un-hybridized orbitals, and are always perpendicular to sigma bonds.
∙ Single bonds are made of a single sigma bond (σ)
∙ Double bonds are made of a sigma bond and a pi bond (σ+π) ∙ Triple bonds are made of a sigma bond and two pi bonds (σ+π+π) o When distributing electrons, sigma bonds get electrons first, then the pi bonds, and then the lone pairs.
NOTE; Hydrogen will never hybridize, because it only has 1 s orbital. Hybridization of Formaldehyde
∙ Formaldehyde has the chemical formula H2CO. Both hydrogens are single bonded (sigma bonded σ) to carbon, and the oxygen is double bonded to carbon (sigma + pi bonds, σ+π) with two left over lone pairs
∙ For its hybridization, the carbon gets the designation sp^2, as it has 3 electron groups (hydrogen, hydrogen, oxygen), and the oxygen also gets the designation sp^2, as it also has 3 electron groups (carbon, lone pair, lone pair).
o It’s important to note that although a double bond consists of 2 electron bonds, the double bond itself only counts as one
Bonding Details of Formaldehyde
∙ Carbon’s bonds; 2 single (σ) bonds and 1 double (σ+π) bond. o That’s a total of 3 sigma (σ) bonds and 1 pi bond (π)
Since pi bonds require unhybridized orbitals and sigma
bonds require hybridized orbitals, that means we need 3
hybrid orbitals for the sigma bonds and 1 normal orbital
for the pi bond.
o Since only the valence electrons hybridize, principal shell 2 is the only one we have to work with (Carbon’s electron
configuration is 1s^22s^22p^2, we get the 2s^2 and the 2p^2) o The valence electrons give us 4 orbitals to work with, the 1 slot in the s orbital and the 3 slots in the p orbital. We need 3 orbitals for hybridization (sp^2) and 1 orbital for non hybridization (the last p orbital). All the orbitals are used and all the electrons are happy.
∙ Oxygen’s bonds; 1 double bond (π+σ) (the other two groups are lone pairs)
o That’s a total of 1 sigma (σ) bond and 1 pi (π) bond, 1 hybridized orbital and 1 unhybridized orbital.
o Because oxygen has those 2 sets of lone pairs for a total of 3 electron groups, it still wants sp^2 orbitals. This means it has
the same hybridization set up as carbon, with the 3 hybridized
orbitals and the one unhybridized orbital.
o Oxygen uses the unhybridized 2p orbital for its pi (π) bond and the first sp^2 orbital for its sigma (σ) bond, and the other two
sp^2 orbitals are taken up by its lone pairs
NOTE; pi (π) bonds CANNOT be rotated. This means that
single bonds are the ONLY bonds that can be rotated, as
both double bonds and triple bonds contain pi bonds.
VERY IMPORTANT: IF YOU’RE STUDYING THIS FROM THE TEXTBOOK, THE TEXTBOOK FALSELY ASSUMES THAT ONLY CENTRAL ATOMS HYBRIDIZE. ALSO, THE DASH IN BETWEEN ORBITALS IS NOT A SUBTRACTION SIGN. DO NOT SUBTRACT THE ORBITALS. Practice example questions 39 and 41 in the textbook, they will be on the test. Answers to them are on blackboard.
Lewis Bond Theory: all about connectivity
VSEPR Theory: all about shapes and geometry
Valence Bond Theory: all about bond formation
∙ Does correctly predict the shape of molecules
∙ But it says orbitals belong to individual atoms
o Remember, diamagnetic v paramagnetic
Diamagnetic means all the electrons are paired up and
the molecule is not attracted to a magnetic field
Paramagnetic means there are unpaired electrons and the
molecule is attracted to a magnetic field
o According to Valence Bond Theory, oxygen should be
diamagnetic, but in experimental practice, that is not the case. Which is why we now have Molecular Orbital Theory. Thanks,
Molecular Orbital Theory
∙ Still combines atomic orbitals in the principal shell, but they form molecular orbitals that belong to the whole molecule. It’s like chemical socialism. This is the biggest difference between Molecular Orbital Theory and Valence Bond Theory.
∙ This has been proved to be the most experimentally correct theory. o Because it’s extremely complicated, if you want to work out the bonds between 3+ atoms, you need a computer program. So, as a result, we will only be working with diatonic molecules
(elements with only 2 atoms)
∙ Orbitals are waves. This means they can be combined constructively or destructively.
o In constructive interference, they waves are in sync and when you put them together they make one super big wave
o In destructive interference, the waves are out of sync and when you put them together, they cancel each other out.
Rules for Combining Atomic Orbitals
∙ Orbitals have to have the same shape. 2 s can combine with 2 s, because they are both spheres. 2 s cannot combine with 2 p, because they are different shapes.
∙ Orbitals have to have the same orientation.
o For the s orbital it doesn’t really matter, because if you turn a sphere upside down and 40 degrees to the left, it’s still in the
same orientation it was before.
o For the p orbitals however, this is a real problem. P orbitals have a px orbital along the x axis, a py orbital along the y axis, and a pz orbital along the z axis.
o So, 2s and 2s are good to go, but 2py and 2pz are not good to go, because even though they’re the same shape they have
∙ When 2 orbitals are in phase (interfering constructively), you get a huge bonding molecular orbital.
∙ When 2 orbitals are out of phase (interfering destructively), you get an anti-bonding molecular orbital that forms a new node.
o Nodes are places where the charge shifts from being positive to being negative, they’re zero spaces.
∙ Constructive Bonding causes the orbital to have a lower bonding energy than the energy of the shells it started with. For example, 1 s orbitals have a higher bonding energy than a 1 s bonding orbital.
∙ Destructive Bonding causes the orbital to have a higher bonding energy than the energy of the shell it started with. For example, a 1 s orbital has a lower bonding energy than a 1 s anti-binding orbital.
When naming molecular orbitals
∙ Sigma (σ) orbitals are on the axis of the bond (s orbitals don’t have axises, they are spheres, so they are always sigma)
∙ Constructively interfering sigma bonds are denoted as σ bonds ∙ Destructively interfering sigma bonds are denoted as σ* bonds ∙ Pi (π) orbitals are perpendicular to the bond axis
o NOTE; which bond axis is THE bond axis is completely arbitrary, people just decided the bond axis should be the x axis, which
means that px actually creates sigma bonds, not pi bonds. Px
and py both behave like they should though.
∙ Constructively interfering pi bonds are denoted as π bonds
∙ Destructively interfering pi bonds are denoted as π* bonds
Orbital Energy Diagrams
∙ When looking at orbital energy diagrams, the original atomic orbitals are to the left and right of the molecular orbitals, which are either a little higher or a little lower than their makers depending on if they’re bonding (lower) or antibonding (higher)
∙ When distributing electrons to the orbitals, rememeber to fill the lowest energy levels first (Aufbau), be sure no two electrons have exactly the same quantum numbers (Pauli) and fill degenerative orbitals each with one electron before doubling up (Hund). Hund’s rule is specifically important, because since σ and σ* have different energy levels, they are not degenerative orbitals. You’d fill σ before filling your 1s, and your 1s before your σ*.
∙ Calculated by subtracting the number of electrons in bonding molecules from the number of electrons in non-bonding molecules, and then dividing by two.
∙ If the bond order is ½ or higher, the molecule exists.
∙ If the bond order is 0 or lower, the molecule does not exist.
∙ The bond order tells you how many bonds to use.
o A bond order of 1 means a single bond
o A bond order of 2 means a double bond
o A bond order of 3 means a triple bond
A ½ bond order means an unpaired electron (a radical)
∙ Bond order predicts bond strength; the higher the bond order the stronger the bonds
∙ Bond order also predicts molecular stability, the higher the bond order the more stable the molecule.