Bonding Flashcards
What are valence electrons?
How would you quickly find the number of valence electrons, by using the periodic table?
Valence electrons are all electrons in the outermost energy level.
Ex: in Nitrogen we have 1s22s22p3, n=2 is the outermost energy level. Add all of those electrons (2 from the s, 3 from the p) to get valence = 5.
Valence electron number can quickly be read off of the periodic table as it will be the column number (counting from the left) that the element is in.
Ex: Nitrogen is in the 5th column from the left, so it must have 5 valence shell electrons.
What are Valence Electrons, and how do they affect an atom’s chemistry?
Valence Electrons are the electrons in the atom’s highest energy subshells. They are chemically relevant because they are the electrons that form chemical bonds.
To find the number of valence electrons, simply count across from the left edge of the Periodic Table to the element in question.
What is the difference in valence electrons for:
Oxygen (O) vs Silicon (Si)?
Oxygen has 6 electrons in the outermost energy level (2 in the 2s and 4 in the 2p).
Silicon has more electrons total, but it only has 4 valence electrons (2 in the 3s and 2 in the 3p).
Define:
The octet rule
The octet rule states that atoms desire eight electrons in their valence shells, as this gives them the electron configuration of a noble gas.
An atom will usually have to bond with other atoms to aquire this electron structure.
According to energy rules, what will cause two atoms to form a bond between them?
Chemical bonds form because they lower the potential energy of the electron clouds. Electrons are shared between atoms across the bond, allowing the final state to be more stable than the two were alone.
In general: bonding proceeds so that as many atoms as possible gain a full octet of valence shell electrons.
What are the three strengths of bond possible between atoms?
Give examples of each.
- single bond: one pair of electrons is shared between two atoms
Ex: O-H bond in H2O, or C-H bond in CH4 - double bond: two pairs of electrons are shared between atoms
Ex: O=O bonds in O2, or C=O bonds in CO2 - triple bond: three pairs of electrons are shared between atoms
Ex: N≡N bonds in N2, or C≡N bonds in CNOH
What are the three noteable exceptions to the octet rule?
- odd-electron species. Molecules which have an odd total number of electrons to distribute (hence some atom will come up lacking)
- incomplete octets. Atoms where attaining a full octet would require too many bonds. H, He, Li, Be, B are all considered incomplete octet species due to the high number of electrons they would need.
- expanded octets. Atoms in period 3 or higher that expand to fill their d-block. The term is a misnomer, as this is no longer an octet, but will have 10, 12, or 14 electrons depending on the central atom being bonded.
Why does Nitrogen not succeed in getting a full octet in Nitric Oxide (NO)?
The most electronegative atom will always gain a full octet first.
Nitric Oxide has 11 valence electrons. Oxygen is more electronegative, so will end up with a full octet (two sets of lone e-pairs and a double bond with N). Nitrogen will be left with 3 lone electrons (and two bonds), as it is less electronegative, for a total of 7 electrons.
Why will it be unlikely for Li to ever attain a full octet?
Lithium has only one valence electron. It would need to acquire 7 additional electrons in order to havea full octet. This is statistically (and structurally) unlikely.
H, He, Li, Be, B are all elements that will from incomplete octets due to the improbability of them acquiring enough electrons.
Why can Phosphorous expand its octet to form 5 bonds to Chlorine in PCl5?
P has expanded all of its electrons into the 3d block, in order to create as many stable bonds with chlorine (and complete chlorine’s octets instead) as possible.
Elements in the third row of the periodic table and beyond often exhibit expanded octets of 10, 12, or 14 electrons. This is due to them expanding into the d-block for greater bonding stability.
Common examples of elements that will expand their octets are P, Cl, S, Xe, and Ar.
What are the three major types of chemical bonds?
- ionic bonds: electrons are transferred from a metal to nonmetal
- covalent bonds: electrons are shared between nonmetals
- metallic bonds: electrons ‘“float” between metals in a lattice
Define and give an example of:
ionic bond
An ionic bond is formed when a metal transfers one or more electrons to a nonmetal. Often on the MCAT this will be column I and II transfering electrons to a halogen.
Ex: The classic example is an “ionic salt” like NaCl. Sodium transfers its electron to chlorine.
How does Coulomb’s law apply to ionic bonds?
Coulomb’s law states that the magnitude of the electrostatic force between two charged particles is directly proportional to the product of the magnitude of each of the charges, and inversely proportional to the square of the distance between the two particles.
This holds for all charged particles, hence can be applied to calculate force between the atoms in an ionic bond.
F ∝ q1*q2 / r2
- q1 and q2: charge magnitude of the atom
- r = distance
If we know that NaCl has a higher Electrostatic force between atoms in its lattice than does KBr, what does that tell us about the strength of the bonds?
NaCl will have stronger bonds between adjacent atoms than KBr, due to the stronger force between atoms.
What would be the result in Force between ions in an ionic compound if we were to 1/2 the distance between adjacent ions?
The Force would be increased by a factor of 4, or quadrupled from the original value.
E ∝ q1*q2 / r2
- q1 and q2: charge magnitude of the atom
- r = distance
Define:
Lattice energy
The lattice energy of an ionic compound is the energy associated with forming a crystalline lattice of the compound from the gaseous ions. This may also be refered to as “heat of formation”.
Note: The value of lattice energy is negative, showing that the formation of an ionic compound is exothermic.
How could we evaluate the Electrostatic energy between ionic bonds?
E ∝ q1*q2 / r
- q1 and q2: charge magnitude of the atom
- r = distance
This holds for all charged particles, hence can be applied to calculate energy between the atoms in an ionic bond.
Note: this value will always be negative, due to the charges always being opposite in an ionic compound.
What will be the result in energy if we double the distance between two atoms in an ionic bond?
The energy will be decreased by a factor of 2, or halved from the initial value. Notice that doubling r in the denominator will effectively be the same as dividing by two.
E ∝ q1*q2 / r
- q1 and q2: charge magnitude of the atom
- r = distance
What type of bond would we expect between atoms in molecules like KBr, CaF2, and LiCl?
These are all ionic compounds, commonly referred to as salts. We expect highly ionic bonds between metals and nonmetals in general.
Classic MCAT examples usually include column I and II bonded with the halogen family.
Define:
covalent bond
A covalent bond forms when a nonmetal bonds with another nonmetal by sharing electrons between them, resulting in an overlap of their electron orbitals.
The image below of methane shows the polar covalent C-H bonds.
What are the three types of covalent bonds?
- polar covalent (electrons are held more closely by the higher electronegative species)
- non-polar covalent bond (electrons are perfectly shared between atoms of the same element)
- coordinate covalent bond (molecules share electrons for greater net stability)
What is the difference in bond strength between molecules in an ionic-bonded compound vs a covalent-bonded compound?
What will this do to melting point and boiling point?
Ionic bonds create stronger intermolecular forces than pure covalent bonds. This is due to ionic bonds being extremely polar (more so than ANY covalent bond, by definition), hence it will attract its substituent molecules more strongly.
The melting point and boiling point will both be higher for the ionic compount, since it will require more energy to pull the molecules apart.
Define and give examples of:
Nonpolar (or pure) covalent bond
In a nonpolar covalent bond, both atoms are the same element hence the electron pair is shared equally. On the MCAT, pure covalent can only be the following diatomics: Br2, I2, N2, Cl2, H2, O2, F2.
(BrINClHOF)
Note: though some Chemistry texts will also include bonds like the C-H bond (since it only measures .4 on the Pauling scale), this is clearly still polar since electrons favor Carbon over Hydrogen.
Define and give examples of:
Polar covalent bond
In a polar covalent bond, the electron pair is pulled closer to the more electronegative atom. The result of this is a bond dipole (one end positive, one end negative), hence where we get the term “polar”.
The atom that is more electronegative has a partial negative charge and the atom that is less electronegative has a partial positive charge.
Any bond between different elements will create some value of bond dipole. While some are rather small dipoles (C-H bond dipole is very slight) the only cases where a bond dipole will NOT form are the diatomics (Br2, I2, N2, Cl2, H2, O2, F2).
Define and give an example of:
Coordinate covalent bond
In a coordinate covalent bond between molecules, one atom from one molecule will contribute both electrons to the bond pair. This creates better stability between both molecules.
Common examples are Lewis Acid/Base pairs: NH3 (N donates the electron pair) and H+ (H+ accepts the electron pair). Nitrogen had a full octet to start with, but was very polar until donating the electrons to this bond; Hydrogen didn’t have any electrons, but now will have a full 1s subshell; both are now more stable.
Lewis Acids and Lewis Bases generally combine to form coordinate-covalent bonded molecules, instead of our common ionic salts, describe specifically why?
A Lewis Acid is a species that will accept an electron pair (ex: Boron in BF3)
A Lewis Base is a species that will donate an electron pair (ex: Nitrogen in NH3)
By definition this is creating a coordinate-covalent bond, and the resluting molecule is called an “adduct” instead of a salt, since it could still separate back off.
Explain why metals, or metallic bonded elements, are good conductors of electricity?
In metallic bonding each metal atom in the crystal structure contributes valence electrons to form a “sea” of delocalized electrons. The “sea” of electrons are able to drift through the electron structure, moving from ion to ion. This movement is what gives these molecules the ability to conduct electricity.
What makes an entire molecule polar?
A molecule is polar when the polar covalent bonds in the molecule add via the molecular geometry to give the entire molecule a positive vs negative end.
Ex: H2O has polar bonds between O-H, with Oxygen more electronegative = – (red in the picture) and Hydrogen = + (blue in the picture). Because both of these bonds point in the same plane (down in the picture) we’ll have a general region of negative (top) and general positive (bottom). H20 must be polar!
What makes a molecule nonpolar?
When a molecule has a balanced charge distribution, such that there is not any one general positive or negative region, it is nonpolar.
Note: Polar covalent bonds can still create a nonpolar molecule if there is symmetry in the geometry, such that the dipoles will cancel and create a balanced molecule.
Ex: in BF3 the B-F bond will pull electrons towards the more electronegative F atoms = – (red in the picture) and away from the B atom = + (blue in the picture). Though these are polar bonds, the trigonal planar symmetry keeps any one side from being net negative vs net positive.
Is CO2 a polar or nonpolar molecule?
CO2 is a nonpolar molecule. Although the C=O bond is polar, the symmetrical shape of the molecule cancels the two dipoles and creates an even distribution of charge.
Oxygen atoms are red, Carbon atom is black.
What is the difference between a bonding pair and a lone pair of electrons?
bonding pair: a pair of electrons that is shared between two atoms across a bond (represented by a line in the Lewis diagram)
lone pair: a pair that is associated with only one atom and stays on just that atom (represented by two close dots in the Lewis diagram)
Define and give the two most important characteristics of::
Lewis structure
Lewis structure is a representation of a molecule that shows how electrons are being shared beween atoms.
Characteristically: In a Lewis structure
1) every line represents a shared electron pair and
2) every dot represents one electron.
Note: in the image below it shows the original atoms on the left, then how confusing it would look if we didn’t have a Lewis structure, then on the far right the proper Lewis structure is displayed.
How would you actually calculate Formal Charge with an easier, shortcut formula?
Formal charge = (# of valence electrons) - (# of lone pair electrons) - (1/2 # of bonding electrons)
In order to calculate formal charge (which accompanies a Lewis structure diagram) we can say:
FC = valence electrons - sticks - dots.
Ex: in CO2, Carbon has 4 valence and 4 sticks, FC=4-4=0. Oxygen has 6 valence and 4 dots and 2 sticks, FC=6-4-2=0.
Define:
resonance structure
A resonance structure is when two or more valid (stable) Lewis structures can be drawn for the same compound.
In resonance structures only the strength of bonds between atoms, not the actual placement of the atoms. If multiple resonance structures exist, the “hybrid” molecule is thought to exist as an average of all of these structures.
Ex: the benzene molecule below has two possible structures, so we consider the C-C or C=C bonds to have an average 1.5 strength no matter what position on the hybrid, due to resonance.
How do you calculate the average bond strength between any two specific atoms in a resonance hybrid molecule?
Average bond strength is quite simply: the average stregths of all bonds possible in that position.
Take any two atoms, count the total number of bonds in that position across all resonance structures, then divide by the total number of structures.
Ex: Nitrate (below), taking the top Oxygen and central Nitrogen, we see that there are 2 bonds + 1 bond + 1 bond =4 total in that position. There are three total structures. Hence the bond strength is 4/3.
Describe the characteristics and give an example of:
Linear molecule
A linear molecule:
- generally has 3 atoms bonded (since any 2 atom molecule is linear by default)
- atoms are arranged 180 degrees apart
- usually contains two double bonds, or one triple and one single bond.
Classic examples include: CO2 (double double) and HCN (single triple)
Describe the characteristics and give an example of:
Bent molecule
A Bent molecule:
- generally has 3 atoms bonded (since there must be one central atom)
- atoms are arranged 104.5 degrees apart
- contains only two single bonds (and two lone e- pairs)
Classic examples include: H2O and CH2
Describe the characteristics and give an example of:
Trigonal Planar molecule
A Trigonal Planar molecule:
- generally has 4 atoms bonded (since we need one central atom and 3 attached)
- atoms are arranged 120 degrees apart
- usually contains single bonds, though double bonds are possible.
Classic examples include: SO3 (double bonds) and BF3 (single bonds)
Describe the characteristics and give an example of:
Trigonal Pyramidal molecule
trigonal pyramidal molecule:
- generally has 4 atoms bonded (1 central and 3 attached below)
- atoms are arranged 107 degrees apart
- contains 3 single bonds, and one lone e- pair
Classic examples include: NH3 and PCl3
Describe the characteristics and give an example of:
Tetrahedral molecule
tetrahedral molecule:
- generally has 5 atoms bonded (1 central and 4 attached below)
- atoms are arranged 109.5 degrees apart
- contains 4 single bonds
Classic examples include: CH4 and CCl4
# Define: Molecular Geometry
Molecular geometry is the actual placement of atoms in a molecule.
Ex: H2O has two Hydrogen atoms placed in a “bent” structure, so this is a bent molecular geometry.
# Define: Electronic Geometry
Electronic Geometry is the position of all bonding electrons and lone e- pairs around one central atom.
Ex: in H2O, Oxygen has two Hydrogens bonded and two lone e- pairs skew to those bonds, for 4 total electron positions in a tetrahedral shape around the Oxygen. H2O has tetrahedral electronic geometry.
Describe an:
sp hybridization
- one s orbital mixes with one p orbital creating two energetically equalivalent hybrid sp orbitals
- the two remaing p orbitals are unchanged and lie at right angles to the plane of the hybrid orbitals
- molecule will be linear in geometry by definition (the hybrid sp orbitals are already linear and only formed in order to bond)
- common examples: C2H2, BeCl2
Describe an:
sp2 hybridization
- one s orbital mixes with two p orbitals.
- the one reamaining unchanged p orbital lies at right angles to the plane of the hybrid orbitals
- resulting molecule will be trigonal planar in geometry
- common examples: H2CO, BF3
Describe an:
sp3 hybridization
- one s orbital mixes with all three of the p orbitals
- resulting molecule will be tetrahedral in geometry
- contains only single bonds on a central carbon atom
- common examples: CH4, CCl4
Define:
sigma bond
A sigma bond occurs when atomic orbitals overlap end to end and result in an accumulation of electron density directly between the nuclei.
Sigma bonds generally exist as single bonds, but a sigma bond makes up the first bond of a double bond or triple bond.
Define:
pi bond
Pi bonds are parallel regions of electron density that overlap and form orbitals alongside an initial sigma orbital. They are generally the result of p orbitals overlaping side by side so the electron density is above and below the internuclear axis.
These make up the second bond in a double bond, and the second and third bonds in a triple bond.
How many sigma and pi bonds does ethene, C2H2, contain?
- The two C-H bonds are single bonds = 2 sigma bonds
- The C≡C form a triple bond = 1 sigma and 2 pi bonds
- Therefore, there are a total of 3 sigma and 2 pi bonds
