Chemical Bonding for MCAT
Lewis structure
representation of a molecule that shows how electrons are being shared between atoms characteristically every line represents a shared electron pair every dot represents one electron.
what makes a molecule non polar
when a molecule has a balanced charge distribution, it is nonpolar there is no comparably positive or negative region in a nonpolar molecule
what makes an entire molecule polar
when the polar covalent bonds in the molecule add via the molecular geometry to give the entire molecule a positive and a negative end. H2O has polar bonds between O-H with oxygen more electronegative. There is a negative region and a positive region. H2O is therefore a polar molecule.
heat of formation or lattice energy is an ------- reaction
exothermic
bent molecule
has 3 atoms bonded has a 104.5 degree bond bond angle contains 2 single bonds and two lone e- pairs H2O SCl2
which compound contains bonds with a higher dipole moment CH4 vs CO2 NaCl vs H2O O2 vs NH3
CO2 because the C+O bonds have a higher dipole moment due to oxygens high electronegativity. C-H bonds are less polar. NaCl has a higher dipole moment than the covalent O-H bonds in H2O polar covalent N-H bonds in NH3 have a higher dipole moment than the pure covalent O_O bond of O2.
is CO2 polar or non polar
CO2 is nonpolar the C=O bond is polar but the symmertical linear shape of the molecule results in cancelled dipoles and an even distribution of charge
electrostatic energy (Ees) of an ionic bond
Ees = q1q2/r is used to calculate energy between the atoms in an ionic bond q1 and q2= charge magnitude of the atom r= distance the value of Ees will always be negative due to the charges always being opposite in an ionic compound
what is the hybridization of carbon in the following H2CO CCl4 CO2 HCN
H2CO- sp2 CCl4- sp3 CO2- sp HCN-sp remember that if x total atoms are bonded to the carbon, then carbons hybridization is sp(x-1) note that sp2 hybridiztion can incolce either the double double bonds in CO2 or the single triple bonds in HCN
why wont Li ever attain a full octet
Li has only one valence electron. It would need to acquire 7 more electrons to fill the octet. this is structurally unlikely. H He Li Be and B are all elements that will form incomplete octets due to the improbability of them acquiring enough electrons
knowing Na and Cl have a larger electronegativity difference than K and Br, what does Coulombs law predict about the strength of the NaCl bond vs that of KBr?
NaCl will have a stronger bond between adjacent atoms than KBr due to stronger electrostatic force between ions.
what is the difference in the valence electron configuration of O and S
O has 6 valence electrons and Si has 4 valence electrons o- 2s2 2p4 Si- 3s2 3p2
why can P expand its octet to form 5 bonds to chlorine in PCl5
P can expand its 5 valence electrons into the 3d block, allowing it to bond to five chlorine atoms and completing those cholorine atoms octets. elements in row 3 and beyond such as P Cl S Xe and Ar often exhibit expanded octets of 10 to 18 electrons. this is due to them expanding into the d block for greater bonding stability
why do central or interior atoms have the highest tendency to hybridize
a central atom is the atom in a molecule that is bonded to more than one other atom, has the highest number of valence electrons and has the lowest electronegativity. all of these qualities make it likely to hybridize in order to have more active valence electroms. central atoms tend to form more bonds too.
according to valence bond theory, what does a chemical bond result from
a chemical bond occurs when one partially filled orbital of one atom overlaps with a partially filled orbital of another atom allowing both to attain a more stable state by sharing electrons. a single bond has one shared pair, a double bond has two shared pairs or 4 total and a triple bond has 3 pairs or 6 total. note that atoms typically form enough bonds to attain a full octet of valence electrons.
why do lewis acids and lewis bases generally combine to form coordinate-covalent bonded molecules instead of common ionic salts
a coordinate covalent bond occurs when one molecule donates both electrons to the covalent bond. lewis acid base pairs combine in exactly that fashion. a lewis acid will accept an electron pair like bonon in BF3 a lewis base will donate an electron pair like N in NH3
which is more electronically repulsive - a pair of electrons in a bond or a non-bonded pair of electrons
a pair of nonbonded electrons has more negative character than a pair of bonding electrons since the nonbonded electrons are not shared between two positive nuclei. so 2 lone pairs will repel each other the most while two bonded pairs will suffer the least repulsion. the repulsion between a lone pair and a bonded pair falls somewhere in between.
dipole moment
any polar bond with a separation of partial positive and negative charge.
octet rule
atoms desire 8 electrons in their valence shells as this gives them a stable configuration like a noble gas
how is the average bond strength between any two specific atoms in a molecule with multiple resonance structures calculated
average bond strength is sumply the average strengths of the bond between the atoms in each resonance structure. in short, total the bonds between the two atoms in each resonance structure, then divide by the total number of structures. ex- to calculate the N-O bond order in nitrate, consider the upper oxygen. there are 2 bonds +1 bond+1 bond= 4 total bonds possible in that position across the 3 unique structures. Hence the average bond strength is 4/3.
difference between a bonding pair and a lone pair of electrons
bonding pair- pair of electrons that is shared between two atoms across a bond represented by a line in a lewis structure. lone pair- pair of electrons that is associated with only one atom and stays on just that atom represented by two close dots on a lewis structure
what energy principle causes two atoms to form a bond between them
bonds form to 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
nonpolar covalent bond
both atoms are of the same element. the electron pair is shared equally. for MCAT, the only pure convalent bonds are the diatomics, Br2, I2, N2, Cl2, H2, O2, F2
how many sigma and pi bonds does ethylene C2H2 contain
c2H2 contains 3 sigma and 2 pi bonds the 2 CH bonds are single bonds meaning there are 2 sigma bonds the C C is a triple bond yielding 1 sigma and 2 pi bonds
what energy principle causes 2 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 are alone. in general bonding proceeds so that as many atoms as possible gain a full octet of valence shell electrons
intermolecular bonds
connect atoms of hydrogen bonds
intramolecular bonds
connect atoms of the same molecule ionic covalent metallic
how does coulumbs law apply to ionic bonds
coulombs law states that the magnitude of the electrostatic force between 2 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 particle. this holds for all charged particles and can be applied to calculate force between atoms in an ionic bond
how is the hybridization of a carbon atom determined
draw the molecule count the number of atoms bonded around the carbon the hybridization of carbon = sp^(x-1) where x is the number of atoms bound to the carbon CH4 has 4 atoms bonded around carbon so its hybridization must be sp ^4-1 = sp3
polar covalent bond
electron pair is pulled closer to the electronegative atom. the result is a bond dipole (one positive end and one negative end- hence polar). the atom that is more elctromegative will carry a partial negative charge and the atom this is less electonegative will carry a partial positive charge. ex- O-H, N-H, C-H
valence electron
electrons in the outermost energy subshells. chemically relevant because they are the electrons that form chemical bonds
how would the force between ions in an ionic compound change if the distance between the adjacent ions was decreased by half
electrostatic force would be 4x the orginal value Fbond q1q2/r^2 Since new R = r/2
resonance structures
exists when 2 or more valid stable lewis structures can be drawn for the same compound. only the strength bonds between atoms varies, not the actual placement of the atoms. if multiple resonance structures exist, the actual molecules structure is an average of all of these. ex- the benzene molecule has two possible structures so the C-C bonds have an average bond order of 1.5 in any given position due to resonance
what is an easier method for calculating formal charge of an atom in a compound, given the compounds lewis structure
formal charge - #valence electons-lines-dots for example, in CO2, carbon has 4 valence electrons, 4 lines, 0 dots. FC = 4-4-0=0. oxygen has 6 valence electrons, 4 dots and 2 lines, so FC= 6-4-2=0
how the formal charge on an atom is calculated
formal charge= (# of valence electrons)-(#of lone pair electrons)-(1/2 # of bonding electrons) formal charge is essentially a fictitious charge assigned to each atom in a lewis structure for the sake of helping distinguish the best lewis structure for a molecule
trigonal pyramidal molecule
generally has 4 atoms (1 central, 3 perpheral) has 107 degree bond bond angle contains 3 single bonds and one lone e-pair NH3 Pcl3
trigonal planar molecule
generally has 4 atoms bonded (one central and 3 peripheral) has 120 degrees bond bond angle classic examples include SO3 (3 double bonds) BF3 (3 single bonds) H2CO (two single bonds, one double bond)
tetrahedal molecule
generally has 5 atoms bonded (1 central 4 peripheral) has 109.5 degree bond bond angles contains 4 single bonds CH4 CCl4
linear molecule
has 3 atoms bonded has atoms arranged 180 degrees apart usually contains 2 double bonds or one tripple and on single bond CO2 double double HCN single triple
ionic bonds has -----melting and boiling pts than covalent bonds
higher more energy required to pull apart the ionic bond which is stronger than the covalent bond
what is a hybrid orbital and what are the 3 most common hybrid orbital types
hybrid orbitals are formed when starndard atomic orbitals combine. they are given a name representing which orbitals have overlapped. the most common hybrid orbitals are sp, sp2, sp3 orbitals sp- an s and a p orbital sp2- s orbital and 2 p sp3- s orbital and 3 p
what type of bond will exist between atoms in molecules like KBr, CaF2, LiCl
ionic bonds these are ionic compounds referred to as salts classic examples of ionic compounds will usually include a cation from Group 1 and II bonded to one or two anions from the halogen family
ionic compound vs covalent compound in terms of bond strength , melting pts and boiling pts
ionic bonds create stronger intermolecular forces than pure covalent bonds. ionic bonds are extremely polar and as such a strong electrostatic interaction exist between the ions. melting and boiling pts will be higher for ionic compounds as more energy is required to pull atoms apart
3 major types of intramolecular bonds
ionic- electrons transferred from a metal to nonmetal covalent- electrons shared between non metals metallic bonds- electrons floar between lattice of metallic nuclei intramolecular bonds connect atoms of the same molecule, unlike intermolecular forces like hydrogen bonds.
lattice energy
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 referred to as heat of formation. the value of lattice energy is negative showing that the formation of an ionic compound is exothermic
bond between metal and nonmetal bond
metal and non metal bonds will be ionic the elctronegative nonmetal withdraws one or more electrons from the electropositive metal leaving them ionically charged.
ionic bond
metal transfers one or more electrons to a nonmetal NaCl Na transfers one valence electron to Cl leaving each with an octet of valence electrons
bonding for metals in their standard states
metals form a lattice of positively charged nuclei in a background sea of free flowing negatively charged electrons. the valance electrons of metals are loosely bound to the atomic cores and can be considered to be effectively unattached
covalent bond
nonmetal bonds with another nonmetal by sharing electrons between them resulting in an overlap of their electron orbital
bond between two nonmetallic elements
nonmetallic elements form covalent bonds nonmetals tend to have similar values of elctronegativity and share electron density fairly evenly when bound together
3 exceptions to the octet rule
odd electron species- molecules like NO must break the octet rule since they have an odd total number of valence electrons to distribute 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 row 3 or higher like P or Cl can hold electrons in their d orbitals. these atoms can hold between 10 and 18 electrons depending on the central atom being bonded.
coordinate covalent bond
one atom from one molecule will contribute both electrons to the bond pair. this creates better stability between both molecules. common example- lewis acid/base pair NH3 (N donates the electron pair) and H+ accepts the electron pair. Nitrogen had a full octet to start with but was very polar donating the electron to the bond with hydrogen. H+ didn't have any electrons but now will have a dull 1s subshell. both are now more stable.
sp hybridization
one s orbital mixes iwth one p orbital creating two energically equivalent hybrid sp orbitals the two remaining unhybridized p orbitals lie at tight angles to the plane of the hybrid orbitals molecule will be linear in geometry C2H2 Be Cl2
sp3 hybridization
one s orbital mixes with 3 p resulting molecule will be tetrahedal four single bonds orented around a central atom usually carbon CH4 CCl4
sp2 hybridization
one s orbital with 2 p orbitals one remaining unhybridized p orbital lies at a right angel to the plane of the 3 hybrid orbitals resulting molecule will be trogonal planar in geometry H2CO BF3
pi bond
pi bonds are parellel regions of electron density that form orbitals alongside an initial sigma bond. they result from p orbitals overlapping side by side so the electron density is above and below the internuclear axis. Pi bonds make up one bond in a double bond and 2 bonds in a triple bond.
when will a molecule with polar bonds still be non polar
polar covalent bonds can still create non polar molecules if the geometry is symmetrical such that the dipoles will cancel and create a balanced molecule example- BF3 , the BF bond will pull electrons toword the more electronegative F atoms. These are polar bonds, but the trigonal symmetry means the molecule has no net dipole.
3 types of covalent bonds
polar covalent- electrons are held more closely by the higher electronegative species nonpolar covalent- electrons are perfectly shared between atoms of the same element coordinate covalent bond- entire molecules share electrons for greater net stability
electronic geometry
postion of all bonding electrons and lone e- pairs around one central atom. H2O has oxygen with two hydrogens bonded and two lone e- pairs in a skew position to those bonds for 4 total electron positions in a tetrahedral shape around oxygen. H2O has tetrahedral electronic geometry while its molecular geometry is bent.
sigma bond
sigma bond occurs when atomic orbitals overlap end to end and result in a accumulation of electron density directly between the nuclei. sima bonds usually exist as single bonds but a sigma bond also makes up one bond of any double or triple bond.
3 types of interatomic bonds
single bond- one pair of electrons shared between 2 atoms like O-H in H2O or C-H in CH4 double bonds- two pairs of electrons are shared between atoms like O=O in O2 or C+O in CO2 triple bond- three pairs of electrons are shared bwtween atoms like N N in N2 or CN in HCN
how many sigma and pi bonds are there in the following single bonds double bonds triple bonds
single bonds- 1 sigma double bond- 1 sigma, 1 pi triple- 1 sigma, 2 pi
ionic bonds are -----than covalent bonds
stronger
molecular geometry
the acutaul placement of atoms in a molecule, ignoring any nonbonded electrons H2O has 2 H atoms in a bent structure so it has a bent molecular geometry
formal charge of an atom in a molecule
the charge it would have if all bonding electrons were shared equally between the bonded atoms. typically a molecules ideal lewis structure will have 0 formal charge on each atom
how would you quickly find the valence electrons by using periodic table
the column number counting from the left represents the elements valence number ex- N is in the 5th column from the left so it has 5 valence electrons. N is 1s2 2s2 2p3 and n=2 is the outermost energy level. add the s2 and p3 electrons to get a valence of 5
how does energy of an ionic bond change if the distance between the two atoms in the bond doubles
the final energy will be 1/2 the initial Eorig=q1q2/R since new R= 2r Enew=q1q2/R q1q2/2r= Eorig/2
why does N not succeed in getting a full octet in nitric oxide (NO)
the most electonegative atom will always gain the full octet first. NO has 11 valence electrons. O is more electronegative so will end up with a full octet (2 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 metals or metallic bonded elements are good conductors of electricity
the sea of electrons is able to drift through the electron structure moving from ion to ion giving these structures the ability to conduct electricity in metallic bonding each metal atom in the crystal structure contributes valence electrons to form a sea of delocalized electrons
valence shell electron pair repulsion theory
the valence shell electron pair repulsion (VSEPR) theory states that electron pairs repel each other. therefore, for a molecule to be at its most stable state, electron pairs should be as far from each other as possible in three dimensional space.
