Chemistry 101 Final Unit 2 Flashcards
What is ionic bonding?
- metals + nonmetals, complete transfer of electrons, forms an extended structures
- IB strength: magnitude of charges + size of ions (ion’s
strength increases as the ion charges increases + ion size
decreases) - ionic solid properties
- hard + rigid, brittle, poor electrical properties
What is lattice energy?
The energy released when gaseous ions combine to form the ionic solid, is proportional to the electrostatic energy (measured strength of IB)
What is metallic bonding?
- metals + metals, detached electrons (sea of electrons) in extended structure
- metallic compounds: malleable/ductile, moderately high mp+bp, good thermal+electrical conductivity
What is covalent bonding?
- nonmetals + nonmetals, typically liquid or gas with low mp+bp, sharing electrons between atoms form molecules
What is bond energy?
Energy required to overcome the attraction, the standard energy change for breaking the bond in 1 mol of gas molecule
- breaking bond: endothermic (BE=+, more Ef than Ei)
What is bond strength + length?
How strong the atoms are joint to each other, depends on mutual attraction between nuclei + shared electrons.
- increase bond order: increase bond strength + decrease
bond length
What is electronegativity?
- atom bonded in molecule’s relative ability to attract shared electrons to itself.
- larger EN: increase in attraction, bond polarity + ionic character (EN difference between cation + anion)
- top-right on period table
- F>O>Cl,N>Br>I,C,S>H (fucking bricks with lisp)
- EN increase across period: if atom’s v-shell less than 1/2 full,
needs less energy to lose e than gain e - EN decrease down group: increased distance between ve +
nucleus (larger atomic radius)
What are homogenous bonds, heterogenous bonds + polar covalent bonds?
- homogenous bonds: nonpolar
- heterogenous bonds: polar
- polar covalent bonds: more EN takes greater share of the bonding e, partial negative charge
Ionic - polar covalent - nonpolar covalent in terms of physical states?
- ionic: solids
- polar covalent: liquids, gases
- nonpolar covalent: gases
What is a dipole moment?
- measure of the separation of + and - electrical charges in molecule (partial charges separated by distance
- electric field: partial charges will align with opposite-
charged electrode
- electric field: partial charges will align with opposite-
- bond has + pole and - pole
- > 2 atom molecules may or may not have overall dipole moment, depends on how the bond dipoles add vectorially
What is resonance + resonance hybrid (true structure)?
- more than one plausible Lewis structure for same molecule: bonding e-density can be detached over more than 2 atoms
- fractional BO happens when there’s partial bonding in resonance hybrid
- total # of bonds divided by # of bond groups
What are the most important Lewis structures?
- complete octets
- low formal charges
- negative formal charges borne by more EN atoms
- separated-like charges
What is formal charge?
- assessing which Lewis structure is the most important contributor, estimating charges on bonded atom
- how it works
- each atom has electrons which “belong to them”
- lone pair: unshared electrons
- bond pair: shared electrons
- formal charge: (ve in free atom) - (assigned e in bonded
atom) - neutral molecule: formal charge sum is 0
- ion: formal charge sum: ion’s charge
- each atom has electrons which “belong to them”
What is the octet rule + its exceptions?
- octet rule: each valence orbital (1 ns + 3 np) can only accommodate 2 electrons (total: 8 electrons)
- exceptions
- odd-electron species: # ve is odd number: place unpaired
electron on least EN atom - incomplete octets: Be, B, Al may have less than an octet
- expanded valence shells: 3rd (or heavier) period elements
may have 10 or 12 electrons around them - any valence electrons remaining after completing octets of
all atoms in a species are given to central atom
- odd-electron species: # ve is odd number: place unpaired
Differences between formal charges and oxidation numbers are…
- oxidation numbers
- doesn’t change from one resonance structure to another
- ionic extreme, electrons given to more EN atom
- formal charges
- changes from one resonance structure to another (# of
bonding + lone pairs can change) - covalent extreme, electrons split evenly
- changes from one resonance structure to another (# of
- neither atom’s oxidation number or formal charge represents actual charge, rather serves to keep track of electrons
- limitations: not useful in choosing correct location
What is VSPER Theory?
- electron pairs arranged around central atom to minimize repulsions
- electron (group) geometry: arrangement of e groups, molecular geometry: 3D structure of atoms in molecule
- no lone pairs: electron + molecular geometry is the same
- groups
- 2 groups: linear (AX2, 180°)
- 3 groups: trigonal planar (AX3, 120°), V-shape (AX2E, <120°)
- 4 groups: tetrahedral (AX4, 105,5°), trigonal pyramidal
(AX3E, <109.5°), V-shape (AX2E2, <109.5°) - 5 groups: trigonal bipyramidal (AX5, 90° + 120°), seesaw
(AX4E, <90° + <120°), T-shape (AX3E2, <90°), linear (AX2E3,
180°)- 90° is axial (vertical), 120° is equatorial (horizontal)
- 2 in axial, 3 in equatorial
- 6 groups: octahedral (AX6, 90°), square pyramidal (AX5E,
<90°), square planar (AX4E2, 90°)- 2 in axial, 4 in equatorial, lone groups on axial line
What is valence bond theory?
A covalent bond forms when orbitals of two atoms overlap, and a pair of electrons are localized in the region between the atoms.
- pair of electrons has opposing spins
- good in phase overlap: strong bonding interactions
- orbitals can hybridize to more appropriate shapes +
orientations to maximize overlap
- hybridization: mathematical mixing of certain
combinations of orbitals, forms new hybrid orbitals
where spatial orientation matches the molecular shapes
we observe
What are the 3 types of covalent bonds in orbital overlaps?
- sigma bond: end-to-end overlap of orbitals, highest e-density along axis of bond
- pi bond: side-to-side overlap of orbitals, e-density above + below axis of bond
- double bond: 1 sigma bond + 1 pi bond: higher e-density
What is the difference between valence bond theory and molecular orbital theory?
VB: pictures a molecule as a group of atoms bonded through overlapping of valence-shell atomic and/or hybrid orbitals occupied by delocalized electrons.
MO: pictures a molecule as a collection of nuclei with orbitals that extend over the whole molecule and are occupied by delocalized electrons.
Molecular orbitals for 2nd period elements…
- pi2pMO - pi2pAMO - sigma2pMO - sigma2pAMO : burger - butterfly wings - candy - 2 drumsticks
- B2, C2, N2: two triangles, O2, F2: elevator buttons
- for bond order
- increase BO: stronger the bond, more stability
- BO > 0: bond exists since molecule is more stable than the
separate atoms - BO = 0: bond doesn’t exist, this is when the molecule is as
stable as the separate atoms
What are the limitations of the localized model? (Lewis structure, VSEPR, VB theory)
Doesn’t explain molecules with odd # of electrons, unpaired electrons and their magnetic properties, and no quantitative BE information
What is HOMO and LUMO?
- HOMO: highest occupied MO
- LUMO: lowest occupied MO
What are the 3 definitions for acids + bases?
- Arrhenius: acids produce H+ and bases produce OH- in aqueous solutions
- acid: H+ source, base: OH- source
- acids + bases react to neutralize each other
- limitations: restricted to reactions in aq solutions, can’t
explain why NH3 is a base
- Bronsted-Lowry: focuses on proton (H+) transfer
- acid: proton donor, base: proton acceptor
- conjugate acid-base pairs related by H+ transfer
- amphoteric: can behave as either acid or base
- Lewis: acid is e-pair acceptor, base is e-pair donor
How do you define acid + base strength?
- acid ionization constant: Ka: produced acid x water / reacting acid, base ionization constant: produced base x OH- / reacting base
- strong acids/bases show complete dissociation into ions
- weak acids/bases show partial dissociation into ions
- since acids + bases come in conjugate pairs: can deduce the relative strengths of the conjugate bases
- ex. if HNO3 is super strong acid, its conjugate base NO3- is
super weak
- ex. if HNO3 is super strong acid, its conjugate base NO3- is
- salts: ionic compounds that result from an acid-base neutralization reaction
What are 2 main factors that affect an acid or base’s strength?
Bond strength + bond polarity
What are binary acids (H-X)?
- acidic proton bonded directly to a non-metal atom X
weaker + more polar the H-X bond: easier to dissociate H+ + the stronger the acid (H-X = H+ + X-) - as you go down the group the H-X bond becomes longer + weaker: increase in acidity (easier to dissociate H+)
- across the period X becomes more EN, thus H-X is more polar: increase in acidity (easier to dissociate H+)