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Year 1 Finals > Chemistry 101 Final Unit 2 > Flashcards

Chemistry 101 Final Unit 2 Flashcards

1
Q

What is ionic bonding?

A
  • 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
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2
Q

What is lattice energy?

A

The energy released when gaseous ions combine to form the ionic solid, is proportional to the electrostatic energy (measured strength of IB)

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3
Q

What is metallic bonding?

A
  • metals + metals, detached electrons (sea of electrons) in extended structure
  • metallic compounds: malleable/ductile, moderately high mp+bp, good thermal+electrical conductivity
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4
Q

What is covalent bonding?

A
  • nonmetals + nonmetals, typically liquid or gas with low mp+bp, sharing electrons between atoms form molecules
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5
Q

What is bond energy?

A

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)

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6
Q

What is bond strength + length?

A

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

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7
Q

What is electronegativity?

A
  • 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)
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8
Q

What are homogenous bonds, heterogenous bonds + polar covalent bonds?

A
  • homogenous bonds: nonpolar
  • heterogenous bonds: polar
  • polar covalent bonds: more EN takes greater share of the bonding e, partial negative charge
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9
Q

Ionic - polar covalent - nonpolar covalent in terms of physical states?

A
  • ionic: solids
  • polar covalent: liquids, gases
  • nonpolar covalent: gases
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10
Q

What is a dipole moment?

A
  • 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
  • bond has + pole and - pole
  • > 2 atom molecules may or may not have overall dipole moment, depends on how the bond dipoles add vectorially
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11
Q

What is resonance + resonance hybrid (true structure)?

A
  • 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
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12
Q

What are the most important Lewis structures?

A
  • complete octets
  • low formal charges
  • negative formal charges borne by more EN atoms
  • separated-like charges
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13
Q

What is formal charge?

A
  • 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
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14
Q

What is the octet rule + its exceptions?

A
  • 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
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15
Q

Differences between formal charges and oxidation numbers are…

A
  1. oxidation numbers
    • doesn’t change from one resonance structure to another
    • ionic extreme, electrons given to more EN atom
  2. formal charges
    • changes from one resonance structure to another (# of
      bonding + lone pairs can change)
    • covalent extreme, electrons split evenly
  3. neither atom’s oxidation number or formal charge represents actual charge, rather serves to keep track of electrons
  4. limitations: not useful in choosing correct location
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16
Q

What is VSPER Theory?

A
  1. electron pairs arranged around central atom to minimize repulsions
  2. electron (group) geometry: arrangement of e groups, molecular geometry: 3D structure of atoms in molecule
    • no lone pairs: electron + molecular geometry is the same
  3. 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
17
Q

What is valence bond theory?

A

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

18
Q

What are the 3 types of covalent bonds in orbital overlaps?

A
  1. sigma bond: end-to-end overlap of orbitals, highest e-density along axis of bond
  2. pi bond: side-to-side overlap of orbitals, e-density above + below axis of bond
  3. double bond: 1 sigma bond + 1 pi bond: higher e-density
19
Q

What is the difference between valence bond theory and molecular orbital theory?

A

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.

20
Q

Molecular orbitals for 2nd period elements…

A
  • 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
21
Q

What are the limitations of the localized model? (Lewis structure, VSEPR, VB theory)

A

Doesn’t explain molecules with odd # of electrons, unpaired electrons and their magnetic properties, and no quantitative BE information

22
Q

What is HOMO and LUMO?

A
  • HOMO: highest occupied MO
  • LUMO: lowest occupied MO
23
Q

What are the 3 definitions for acids + bases?

A
  1. 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
  2. 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
  3. Lewis: acid is e-pair acceptor, base is e-pair donor
24
Q

How do you define acid + base strength?

A
  • 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
  • salts: ionic compounds that result from an acid-base neutralization reaction
25
Q

What are 2 main factors that affect an acid or base’s strength?

A

Bond strength + bond polarity

26
Q

What are binary acids (H-X)?

A
  • 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+)

Year 1 Finals (4 decks)

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