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Unit #2 Flashcards

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1
Q

what is the average trend in melting and boiling point

A

down a group, they decrease
from left to right, they increase until the metalloids then decrease significantly

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2
Q

what is the trend in physical properties

A

they are very difficult to predict across a period (mp, hardness and conductivity) but usually increase until metaloids then tend to decrease

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3
Q

why are metalloids “hard solids”

A

it is due to the strong covalent bonds holding atoms together in their extensive covalent network

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4
Q

what is the 2 important classes of reactions?

A

reduction-oxidation (redox) and acid-base reactions

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5
Q

what are redox reactions

A

the tendency of a species to lose or gain electrons, there is a transfer of electrons (reflected in a change of oxidation state)
- RA : loses e-s, is oxidized by OA, oxi number increases
- OA : gains e-s, is reduced by RA, oxi number decreases
steps:
- write rxn
- determine 1/2 rxns (which is being oxidized and reduced)

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6
Q

identify whether these are good oxidizing or reducing: Mg, O2, MnO4-

A

Mg - RA (goes to 2+)
O2 - OA (goes to 2-)
MnO4 - OA (@ 7+ so it wants to go down)

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7
Q

what are acid-base reactiosn

A

metal oxides and nonmetal oxides differ in their reactivity w/ water
- basic oxide: Li2O + H2O -> 2LiOH (metal)
- acidic oxide: SO2 + H2O -> H2SO3 (non-metal)
aka anhydrides (w/o water)
periodic trend: basic oxides -> acidic oxides
metalloids are usually amphoteric

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8
Q

what are the results of these reactions?
- MgO + H2O
- SO3 + H2O

A
  • MgO + H2O -> Mg(OH)2
  • SO3 + H2O -> H2SO4
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9
Q

why do bonds form?

A

electrostatic attractions: bonding lowers the potential energy between positive and negative particles -> lower energy = more stability
quantum mechanics perspective: atoms form bonds in order to make their outer shell more stable

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10
Q

what properties play a role in bonding

A

properties such as the IE, EA, Zeff, and size all play a role in their bonding (how they come together and bond).
metallic behaviour (Larger size, low IE) increases down a ground and decrease left to right across a period

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11
Q

what are the 3 types of bonding

A

ionic bonding, covalent bonding, and metallic bonding

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12
Q

what is ionic bonding

A

metals and nonmetals bonding with a complete electron transfer (forms cations and anions) arranged in a extended structure.
formula = lowest whole number ration
strength (related to coulombs law, E ∝ q1q2/d) is related to the magnitude of charges and the size of ions (lattice energy is released when the gaseous ions combine to form the ionic solid)

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13
Q

what physical properties does the ionic bonding model account for

A
  • hard and rigid - the positive and negative ions are strongly attracted to eat other and are difficult to separate (req lots of energy)
  • brittle - when stress is applied to the ionic lattice, the layers shift slighty -> ions of like charges are force closer together, this increases electrostatic repulsion which results in the structure breaking down
  • poor conductivity in solids but good in liquid - solid ionic compounds do not conduct because the ions are locked to a rigid lattice. the dissociation of ions in the liquid phase allow for the ions to move out of the lattic structure and conduct electricity
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14
Q

what is covalent bonding

A

nonmetal and nonmetal sharing electrongs between atoms (form a molecule). each atom holds onto its own electrons tightly (high IE) and attracts other electrons as well (highly negative EA)
- typically exist as liquids or gases w/ low mp and bp. to explain you need to consider the distinction between the strong covalent bonds within molecules and the weak forces between molecules

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15
Q

what is metallic bonding?

A

metal and metal, delocalized “sea of electrons” in an extended structure.
- typically malleable and ductile w/ moderately high mp and bp, as well as excellent conductors of heat and electricity

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16
Q

what physical properties does the metallic bonding account for

A
  • malleability and ductility - can be deformed because of the electron sea prevents repulsions among the cations
  • moderately high mp: cations can move without disrupting their attraction to surrounding electrons
  • thermal/electrical conductivity: mobile, delocalized electrons in the metallic structure
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17
Q

what is bond energy

A

the energy required to overcome the attraction and is defines as the standard enthalpy change for breaking the bond in 1 mol of gaseous molecules
- breaking a bond is always endothermic (req energy)
- forming a bond is always exothermic (release energy)

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18
Q

what is bond strength related to

A

the strength of the bond depends on the magnitude of the mutual attraction between bonded nuclei and shared electrons
- be energy increases with increasing the bond orders and decreasing the bond length
- the more bonds there are, the stronger the bond is

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19
Q

rank the relative bond length and strengths and explain why: Si-F, Si-C, Si-O

A

Si-F>Si-O>Si-C (strength) - Si-F<Si-O<Si-C (length)
Atomic radius decreases across a period so flourine is smaller than O which is smaller than C. Therefore, bond length decreases across a period going C->O->F and bond strength increases

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20
Q

as bond order increases, bond length
as bond order increases, bond energy

A

decreases
increases

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21
Q

the net energy change results:

A

from the difference in bond energies:
ΔHrxn » Sum of BE (bonds broken) – Sum of BE (bonds formed)
exo: energy released during the formation of bonds is greater than the energy required to break bonds
endo: energy released during bond formation is smaller than the energy required to break bonds
In other works: the heat released or absorbed during a chemical change is due to differences between reactant bond energies and product bond energies

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22
Q

what is electronegativity?

A

the relative ability of an atom bonded within a molecule to attract shared electrons to itself. the larger the number, the stronger the attraction of the electrons by that atom. It is a relative scale in which atoms are ranked on their ability to attract electrons
- homogeneous bonds -> non-polar
- heterogenous bonds -> polar (the more EN atom takes a greater share of the bonding e- = partial negative charge while the other is slightly positive)

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23
Q

what are the electronegativity trends?

A

REMEMBER F>O>Cl,N>Br>I, C, S>H
- across a period electronegativity increases - if the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one
- down a group electronegativity decreases - there is an increased distance between the valence electrons and nucleus (larger AR)

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24
Q

increase in ΔEN results in:

A

larger partial charges and higher partial ionic character. decreasing ΔEN the bond becomes more covalent and the character of the substance

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25
Q

rank the bond polarity in these substances: SCl2, PCl3, SiCl4

A

S-Cl < P-Cl < Si-Cl *polarity depending on each bond, not the net polarity

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26
Q

what are dipole moments

A

molecules with partial charges separated by a distance d, possess a dipole moment. molecules with dipole moments are aligned in an electric field
- when an electric field is applied, the partial charges will align w/ the oppositely charged electrode (asymmetrical distribution of electrons)
- a bond contains a positive end and negative end (together they are termed the dipole)
- need to consider all bonds and geometry

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27
Q

whats the basic overview of creating lewis dot structures

A
  1. determine the total # of valence elctrons
  2. determine how atoms are connected (central atom is usually the LEAST EN atom, H and F are always terminal and only form 1 bond)
  3. draw a skeletal structure by joining atoms w/ single bonds (-2 electrons/bond)
  4. distribute remaining e- pairs (complete octets of terminal atoms, distribute remaining electrons around central atoms)
    WHEN THINKING OF CONNECTIVITY, THE WAY THAT IT IS WRITTEN WILL USUALLY GIVE YOU HINTS
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28
Q

what is resonance

A

when more than one plausible lewis structure can be drawn. the true structure is a resonance hybrid of the contributing lewis structures. the concept of resonance accounts for the fact that bonding electron density can be DELOCALIZED over more than two atoms. on exams draw the diff. structures not resonance
- leads to fractional bond orders

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29
Q

the most important lewis structures have:

A
  • complete octets
  • low formal charge (+ or -) but ideally zero
  • negative formal charges borne by more electronegative atoms
  • separated like charge (same nonzero charges on adjacent atoms is not preferred)
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30
Q

what is the formula for fractional bond order?

A

e- pairs/bonded atom pairs

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31
Q

what is a basic overview of determining formal charge

A
  1. each atom is assigned electrons which belong to them (lone pair electrons belong entirely to the one atom, bond pair electrons are split b/w two atoms)
  2. formal charge = valence e of free atom - (lone pair e- +1/2 bond pair e)
  3. for a neutral molecule, sum of formal charges is zero, for an ion, the sum of the formal charges equals the charge on the ion
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32
Q

octet rule exceptions

A
  1. odd-electron species (place unpaired e- on least electronegative atom (ex. NO))
  2. incomplete octets (Be, B, Al may have less than an octet (ex. BF3))
  3. expanded valence shells (3rd period of heavier elements may have 10 or 12 electrons around them (ex. ClF3))
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33
Q

what is the difference between formal charges and oxidation number?

A

oxidation numbers: do not change from one resonance structure to another (EN of the atoms do not change)
- ionic extreme, electrons assigned to more EN atom
formal charges: do change -> b/c the #s of bonding and lone pairs change
- covalent extreme, electrons split evenly
NEITHER AN ATOMS FROMAL CHARGE NOT ITS OXIDATION NUMBER REPRESENTS AN ACTUAL CHARGE ON THE ATOM; BOTH OF THESE TYPES OF NUMBERS SERVE SIMPLY TO KEEP TRACK OF ELECTRONS

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34
Q

what are the two main concepts of vsepr

A
  1. valence electrons are present either as bonded electrons localized between atoms or as lone pairs an atom
  2. lone pair electrons and bonded electrons try to minimize like-charge repulsions by arranging themselves as far apart as possible
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35
Q

why do we need effective pairs

A
  • according to vsepr theory, bonded electrons are located directly between the 2 atoms, therefore it makes no difference to the repulsions if there are 1, 2, or 3 pairs as they are all effectively in the same region
  • lone pairs are located around 1 atom, theyre less constricted in their movements and therefore we could their contributions to repulsions separately
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36
Q

what is the ideal geometry of 2 effective pairs

A

central atom (A) is bonded to 2 atoms (X), minimal repulsions achieved at 180 degrees, the geometry is called linear

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37
Q

what is the ideal geometry of 3 effective pairs

A

central atom (A) is bonded to 3 atoms (X), minimal repulsions achieved @ 120 degrees, the geometry is called trigonal planar (3 terminal atoms form a triangle w/ all atoms in the same plane)

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38
Q

what is the ideal geometry of 4 effective pairs

A

central atom (A) is bonded to 4 atoms (X), minimal repulsions achieves @ 109.5 degrees, the geometry is called tetrahedral

39
Q

what is the ideal geometry of 5 effective pairs

A

central atom (A) is bonded to 5 atoms (X), minimal repulsions achieved in 2 planes, axial and equatorial
- equatorial, 3 eps arranged @ 120 degrees on plane
- axial, 2 eps arranged perpendicular to equatorial, @ 180 degrees
bond angle between axial and equatorial atoms (90 degrees) is also reported
the geometry is called trigonal bipyramid

40
Q

what is the ideal geometry of 6 effective pairs

A

central atom (A) is bonded to 6 atoms (X), minimal repulsions achieves @ 90 degrees, the geometry is called octahederal

41
Q

what are the nonideal geometries of 3 effective pairs

A

AX2E (CA, 2X, 1 lone pair). bond angle <120 degrees, the geometry is known as bent or v-shaped
*smaller because the lone pair is located on 1 atom only rather than bonded pairs between 2. this = more freedom and causes stronger repulsions. to compensate, the bonded pairs reduce the distance between them (narrow the angle)

42
Q

what are the nonideal geometries of 4 effective pairs

A
  1. AX3E -> trigonal pyramid
  2. AX2E2 -> bent or v-shaped
    *don’t need to memorize angles but need to recognize that for molecules with the same numbers of effective pairs, the bond angle decreases w/ the increase of lone pairs
43
Q

what are the nonideal geometries of 5 effective pairs and which plane will the lone pairs be in?

A
  1. AX4E (LP = equatorial) -> seesaw
  2. AX3E2 (LP = equatorial) -> T-shaped
  3. AX2E3 (LP = equatorial) -> linear

strongest repulsions in the X5 occur between axial-equatorial effective pairs w/ 3 bond angles of 90 degrees. If the lone pair in AX4E is located in an axial position, 3 90 degree repulsions will remain. If the lone pair in AX4E is located in an equatorial position, 2 90 degree repulsions will remain therefore the lone pair is likely to be in the equatorial position
*don’t need to memorize angles but need to recognize that for molecules with the same numbers of effective pairs, the bond angle decreases w/ the increase of lone pairs

44
Q

what are the nonideal geometries of 6 effective pairs?

A

all AX6 are equal therefore lone pairs can replace any bond, resulting geometry is a distorted square pyramid with bond angle <90
1. AX5E -> square pyramid
2. AX4E2 (min repulsion achieved when LP are opposite each other) -> square planar

45
Q

what is a general overview of how you draw vsepr diagrams

A
  1. draw a plausible lewis structure
  2. establish the electron-group arrangment of AXmEn
  3. determine the molecular shape
  4. predict any deviations from ideal angles
    - repulsive forces decrease in order: lone pair-lone pair > lone pair-bond pair > bond pair-bond pair
  5. is the molecule polar or nonpolar (take vector sum)
46
Q

how do you determine the vsepr of complex molecules

A

you have to determine the geometry around each central atom.
ex. CH3COOH, have to determine the geometry around C1, C2, and O

47
Q

how can we explain molecular shapes based on the interactions of atomic orbitals

A

via valence bond theory
the basic principle of VB theory is that a covalent bond forms when oprbitals of two atoms overlap, and a pair of electrons is localized in the region between the atom.
quantum mechanics perspective: overlap of the two orbitals means their wave functions are in phase -> constructive interference

48
Q

what are the 3 principles of valence bond theory

A
  1. the pair of electrons have opposing spins (exclusion principle: the space formed by overlapping orbitals has a maximum capacity of 2 electrons with opposite (paired) spins)
  2. good, in-phase overlap leads to strong bonding interactions (the greater the overlap, the closer the nuclei are to the free electrons and the stronger the bond
  3. orbitals can hybridize to more appropriate shapes and orientations, to maximize overlap (the mathematical mixing of certain combinations of orbitals)
49
Q

what are features of hybrid orbitals and what are they

A

they are new orbitals whose spatial orientation match the molecular shapes we observe
features:
- the number of hybrid orbitals formed = the number of atomic orbitals
- the type of hybrid orbitals formed varies with the types of atomic orbitals mixed
- the shape and orientation of a hybrid orbitals maximize its overlap with the orbital of the other atom in the bond

50
Q

is hybridization a physical process

A

no, it is a hypothetical process that is invoked to explain molecular shapes that are observed experimentally

51
Q

what’s the hybridization process of electron groups if 2, 3, 4, 5, and 6

A

2 - mixes one s and one p to form two sp orbitals w/ a linear spatial arrangement and leaving 2 p unhybridized
3 - mixes one s and two p to form three sp2 orbitals w/ a trigonal planar and leaving 1 p unhybridized
4 - mixes one s and three p to form four sp3 orbitals w/ a tetrahedral and leaving nothing unhybridized
5 - mixes one s, three p, and one d to form 5 sp3d orbitals w/ a trigonal bipyramidal shaped and 4 d unhybridized
6 - mixes one s, three p, and two d to form 6 sp3d2 orbitals w/ a octahedral shape and 3 p unhybridized

52
Q

what are the sp hybridization shape

A

LOOK AT NOTES
- it looks like a large hourglass w/ one side massive and the other mini

53
Q

what are the sp2 hybridization shapes

A

LOOK AT NOTES
- trigonal planar, other molecules have a end-to-end connection (hourglass)

54
Q

what are the sp3 hybridization shapes

A

LOOK AT NOTES
- tetrahedral shape, typically end-to-end connection

55
Q

what are the sp3d hybridization shapes

A

LOOK AT NOTES
-trigonal bipyramidal, other molecules end-to-end to these orbitals.
*lone pair placements are the same as if they would be in the vsepr diagrams, just don’t have anything that bonds to it

56
Q

what are the sp3d2 hybridization shapes

A

LOOK AT NOTES
- octahedral while other molecules hybridizing end-to-end

57
Q

how do you create bonding using a hybridization scheme w/ valence bond theory

A
  1. write a plausible lewis structure
  2. from vsepr, predict the e- group arrangement
  3. choose a hybridization scheme for the central atom based on the electron groups arrangement (repeat for terminal atoms)
58
Q

what are the limitations of hybridization

A

vsepr theory would predict bond angles slightly below 109.5 (like water) and valence bonding theory would yield sp3 hybridization for the central atom (does not account for bond angles in large nonmetal hydrides)

59
Q

what are sigma bonds

A

end-to-end overlap of orbitals, with electron density along axis of bond
- shaped like an ellipse rotated about its long axis
- all single bonds
- is always the result of the overlap of hybrid orbitals

60
Q

what are pi bonds

A

side-to-side overlap of orbitals, with electron density above and below axis of bond
- has two lobes of electron density, one above and one below the sigma bond axis. the two electrons in one ok bond occupy both lines
- always the result of the overlap of unused p orbitals or unused d orbitals for expanded octets

61
Q

what kind of bonds do double bonds consist of

A

one sigma bond and one pi bond which increases electron density between the nuclei

62
Q

go review 6.2 orbital overlap and types of covalent bonds

A

no srsly go

63
Q

what is molecular orbital theory

A

quantum mechanics are now applied to molecules. LDS give structure, VSEPR gives shapes and hybridization give the qualitative bonding. in all of these models electrons are localized b/w atoms

64
Q

what are the distinctions between valence bond and molecular orbital theories

A
  • VB pictures a molecule as a group of atoms bonded through overlapping of valence-shell atomic and/or hybrid orbitals occupied by localized electrons
  • MO theory pictures a molecule as a collection of nuclei with orbitals that extend over the whole molecule and are occupied by delocalized electrons
65
Q

what is the linear combination of atomic orbitals

A

you can either add the wave functions or subtract them
- adding them together reinforces each other forming a bonding MO with the electron density b/w the nuclei
- subtracting them cancels each other out forms an anti-bonding MO which forms a node (0 density) between nuclei

66
Q

what is H2 in bonding MOs

A

we construct MOs as a linear combination of the 1s atomic orbitals on each H atom to form two new MOs
AO of HA + AO of HB = bonding MO of H2
AO of HA - AO of HB = antibonding MO of H2
- bonding MO (one large lobe) has less energy than antibonding MO (2 balls w/ a node)
- bond order = 1/2(2-0) = 1

67
Q

what is the relation between atomic orbitals and its molecular orbitals for sigma MOs

A
  • a bonding MO is lower in energy that the atomic orbitals that form it. the electron density is spread mostly between the nuclei and nuclear repulsions decrease and the nuclei electron attraction increase
  • two electrons in this MO can delocalize their charges over a large volume that they could in nearby, separate atomic orbitals (lowers electron repulsions)
  • an antibonding MO is high in energy than the atomic orbitals that form it. with most of the electron density the internuclear region and a node between nuclei, nuclear repulsions increase
68
Q

What are 4 features of MO diagrams

A
  1. the number of MOs equals the number of atomic orbitals
  2. combining two atomic orbitals gives a bonding MO and an antibonding MO
  3. In ground state configurations, electrons occupy lowest energy MOs upwards while respecting Pauli’s and Hund’s principles
  4. bond strength can be measure by the calculated bond order (bond order = 1/2(# of bonding e-s - # of antibonding e-s)
    *the MO diagram shows the relative energy and number of electrons for each MO, as well as for the atomic orbitals from which they are found
69
Q

why is He2 an unstable species

A

both the bonding and antibonding orbitals are filled
- stabilization from the electron pair in the bonding MO is cancelled by the destabilization from the electron pair by antibonding MO
- bond order: 1/2 (2-2) = 0, since the bond order is zero, this species is not stable, molecule likely does not exist
***the lower the energy, the more stable a species is

70
Q

what is the generality w/ bond order and stability

A

bond order > 0: the molecule is more stable than the separate atoms, so it will form. (H2, bond order is 1)
bond order = 0: the molecule is as stable as the separate atoms, so it will not form (occurs when = numbers of electrons occupy bonding and antibonding MOs)
*the higher the bond order the more stable the molecule is

71
Q

how do 2s and 2p orbitals combine

A
  • 2s combine to form sigma bonding and antibonding MOs
  • 2p combine to form sigma bonding and antibonding MOs and pi bonding and antibonding MOs
72
Q

explain why the splitting of energy between bonding and antibonding MOs is greater for sigma bonding and antibonding vs pi bonding and antibonding

A

atomic p orbitals overlap (combine) more extensively end to end than side to side therefore the sigma MO is usually lower than the pi MO (large stabilization w/ sigma). we also find that the destabilizing effect of the sigma antibonding MO is greater than the pi antibonding MO.
with the increase in # of nodes, the energy of the MOs increase as well (destabilization)

73
Q

how are sigma and pi MOs different in early diatomic molecules of early elements

A

the MOs are reversed in energy in this period, whereas they have the expected energy ordering for the diatomic molecules of the later elements.
*need to know these two different splitting patters where the pi2p MO is either lower in energy than the sigma2p (B, C, N) or higher than the sigma 2p (O, F, Ne)

74
Q

if C2 is found to be diamagnetic, what is it’s ground-state electron configuration

A

C2: (σ2s)^2(σ2s*)^2(π2p)^4

75
Q

what are limitations of the localized model

A

it does not explain molecules with
- odd #s of electrons (ex. NO)
- unpaired electrons and their magnetic properties
- no quantitative bond energy information

76
Q

what are delocalized electrons w/ the bond MO theory

A

VB theory treats all bonds as localized, whereas MO theory treats all bonds as completely delocalized.

77
Q

how do we use the MO model w/ excited electrons

A

for ground state molecules, the e- fill lower MOs first
*note that an e- can be promoted to a higher energy Mo creating an excited state of the molecule
* denotes a excited state
usually, excitation between the HOMO and LUMO has the highest probability of occurring. it becomes less stable b/c the bond order decreases by going to an anti-bonding orbital

78
Q

overall summary of molecular orbital theory w/ uses and drawbacks

A

summary: atomic orbitals combine to form bonding and antibonding molecular orbitals that extend over the entire molecule (e-s delocalized). it describes the bonding and antibonding interactions based on which orbitals are filled.
uses: predicts the arrangements of electrons and their energy levels in molecules. explains paramagnetism/diamagnetism and makes it possible to predict the properties of hypothetical molecules and ions
drawbacks: easily used for diatomic molecules of the first- and second- period atoms; becomes more comples as the size of the molecule increases

79
Q

why is the concept of acids and bases important

A

it is an important organizing principle of many reactions. these species also appear in many aspects of our lives (pH)

80
Q

how can we identify acids and bases (4)

A

definitions, measure pH, Ka values, and molecular structure/periodic trends

81
Q

what are the different definitions of acids and bases (3)

A
  1. arrhenius
  2. Brøsted-Lowry
  3. lewis
82
Q

what is arrhenius’ definition of acids and bases

A

acids produce H+ and bases produce OH- in aqueous solutions
acid: source of H+(aq)
base: source of OH-(aq)
acids and bases react to neutralize each other - acid + base -> salt and water
limitations, restricted to reactions in aqueous solutions, cannot explain why certain species are acids or bases

83
Q

what is bronsted-lowry’s definition of acids and bases

A

acid is the proton donor to base’s proton acceptor.
conjugate acid-base pairs are related by transfer of H+ (a substance can be amphoteric depending on what other substances are present

84
Q

what is lewis’ definition of acids and bases

A

THE DONATION AND ACCEPTANCE OF AN ELECTRON PAIR TO FORM A COVALENT BOND IN AN ADDUCT
acid: electron-pair acceptor to form a bond (must have a vacant orbital)
base: electron-pair donor to form a bond (must have a lone pair to donate)
requires that a base have an electron pair to donate, so it does not expand the classes of bases, however, it greatly expands the classes of acids.
the product of a lewis acid-base reaction is an adduct, a single species that contains a NEW covalent bond

85
Q

how do we quantify the strength of a acid or a base

A

by relating it to the magnitude of the equilibrium constant K
- Ka = the acid ionization constant
- Kb = the base ionization constant
strong acids/bases, complete dissociation into ion
weak acids/bases, partial dissociation into ions

86
Q

what are the strong acids and bases

A

acids: HNO3, HCl, HBr, HI, H2SO4, HClO4
bases: group 1 and 2 metals w/ OH-

87
Q

how can we deduce the relative strength of conjugate pairs

A

strength of conjugate base increases while the strength of acids decrease
- anion from a strong acid = no acid-base properties
- cation from a strong base = no acid-base properties
- anion from a weak acid = weak base
- cation from a weak base = weak acid

88
Q

what factors affect the strength of an acid or a base

A

2 main factors: bond strength and bond polarity

89
Q

what are binary acids

A

H - X
the acidic proton is bonded directly to a nonmetal atom X. the weaker and more polar the H - X bond, the easier it is to dissociate H+ and the stronger the acid.
a) down a group, acidity increases -> bond length and atom size increase
- the bond becomes longer and weaker so it is easier to dissociate H+ (why HI is the stronger acid), easier to break the H-X bond
b) across a period, it increases as it goes left to right
- going to the right, X becomes more electronegative, so the H-X bond becomes more polar (tending to have greater partial positive charge on H and partial negative charge on X) and H+ is more liekly to dissociateR

90
Q

Rank in order of base strength: OH- vs SH-

A

OH- is a stronger base because H2S is more acidic than H2O therefore the conjugare of H2S is a weaker base

91
Q

what are oxiacids

A

H - O - X
- the acidic proton is bonded to an O atom, which in turn is bonded to another element X (H IS ALWAYS BONDED THROUGH O)
- a more polar O - H bond leads to a stronger acid because it is easier to dissociate H+
a) high electronegativity of X
b) larger number of terminal O atoms

92
Q

rate strength of acids: HOI, HOCl, HOBr

A

HOI < HOBr < HOCl
- Cl- is more electronegative than Br and I (respectively), the electron density shifts away from H-o, resulting in HOCL being the easiest to dissociate H+ therefore making it the strongest acid

93
Q

rate strength of acids: HOCl, HClO2, HClO3, HClO4

A

HOCl < HClO2 < HClO3 < HClO4
- extra oxygens -> the higher the electron density is withdrawing power, electron density shifts away from the H-O bond therefore stronger the acid
- highly electronegative O atoms withdraw electron density away from the O-H bond, so the more there are, the stronger the acid

94
Q

what makes a good RA and a good OA

A

for metals - reactivity increases down a group LOW IE, SMALL EA = GOOD RA (low oxi states)
for nonmetals - reactivity decreases down a group, LARGE IE, LOW EA = GOOD OA (high oxi states)

Chem 101 (4 decks)

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