Brainscape's Knowledge GenomeTM


Top Flashcards (Certified)
All Flashcards

Tufts - Chem 1 > Exam 2 > Flashcards

Exam 2 Flashcards

(87 cards)

Start Studying
1
Q

Lewis structure rules

A
  1. Atom lowest electronegativity in center (except for H cannot be placed as center)
  2. Find total number of v.e
    (Add +1 v.e for negative charge, subtract -1 v.e for positive charge)
  3. Connect all atoms using single bonds (lines), add remaining electrons to terminal atoms (pairs)
  4. If central atom has < 8 electrons, move lone pair from outside atom to make multiple bonds
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

lewis exceptions

A
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Resonance structures:

A

same arrangement of atoms but diff placement of electrons

distribution of electrons is an avg of all lewis structures

Resonance hybrid: superposition/avg of two resonance structures

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

How do lewis structures depict electrons incorrectly?

A

depict electrons as localized

In nature, electrons are delocalized: density spread over entire molecule

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Bond order

A

(# bonding e - # antibonding e) / 2

Higher bond order = stronger bond = shorter bond length = higher bond energy

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Equivalent vs. non-equivalent resonance structure:

A

Equivalent resonance structure: same distribution formal charges

Non-equivalent resonance structure: different distribution FC
(Don’t equally contribute to resonance hybrid)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Formal charge (FC):

A

FC = #v.e periodic table – #v.e in bonded atom

hypothetical charge an atom would have if we could redistribute the electrons in the bonds evenly between the atoms

Sum formal charges on all atoms = overall; charge molecule

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Criteria for choosing greatest contribution to resonance hybrid:

A
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

What are some drawbacks of lewis structures?

A

doesn’t explain shape/geometry of molecule (effect properties)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Valence Shell Electron Pair Repulsion Theory (VSEPR)

A

predict shape/geometry molecule from lewis dot structure

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Molecular shape notation

A
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Electron pair geometry:

A

all regions where electrons, bonds, lone pairs are located

Same as molecular structure when there are no lone pairs around central atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Molecular structure:

A

location of the atoms, not electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Valence bond theory (VB):

A

When atoms approach each other for bonding, atomic orbitals of one atom overlap with the atomic orbitals of the other atom

Each of the atomic orbitals that overlap should have 1 electron with opposite spin

After orbitals overlap, pair of electrons occupy the overlapped region

Strength bond: greater orbital overlap (closer nuclei is to bonded electrons) = stronger bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Hybridization:

A

Hybridization: orbital mixing to form hybrid orbitals (linear combination of atomic orbitals)

  1. Number hybrid orbitals formed = number of atomic orbitals combined
  2. Hybrid orbitals formed are equivalent in shape and energy
  3. Hybrid orbitals are more effective in forming bonds than unhybridized orbitals
  4. Hybrid orbitals orient themselves in 3D to max distance between then and min repulsions between electrons
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Hybridized vs non-hybridized orbitals

A

Non hybridized orbitals are non equivalent in energy and shape, whereas hybridized orbitals are equivalent in shape and energy

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

rule for number of hybrid orbitals?

A

Number of electron groups around central atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Sigma vs. pi bond

A

Sigma bonds (σ): electron pair is in the region centered on the internuclear axis
- Lobes point toward each other (end to end overlap)
- Free rotation around sigma bond

Pi bonds (π): electron pair in the region above and below the internuclear axis
- Orbitals that are parallel to each other (sideways overlap)
- Restricted rotation around pi bond (transiomer)

π bond weaker than σ bond bc sideways overlap is less effective

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Polar vs. nonpolar

A

POLAR:
- asymmetric electron distribution around molecule
- Dipole moment: measure polarity
- Arrow pointing toward more electronegative

NONPOLAR:
- symmetric electron distribution, electronegativity difference = 0

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

In a molecule with 2+ atoms, what does the overall polarity depend on?

A

Individual bond polarity (polar bonds: asymmetry, individual bond dipoles dont cancel: asymmetry in molecular shape)

Shape of molecule

Effect of molecular polarity on behavior:
Boiling point (intermolecular forces)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Intermolecular forces (IMF):

A

Intermolecular forces: attraction between molecules with partial charges, or between ions/molecules

Weaker than bonding forces (F = q1q2 / r^2)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

IMF: Dipole dipole forces

A

ALL molecule with net dipole moment have dipole dipole forces

  • Positive pole of one polar molecule attracts the negative pole of another
  • Same molar mass: ↑ dipole moment = ↑dipole dipole forces = ↑higher boiling point
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

IMF: hydrogen bond

A

extreme form of dipole dipole forces

Hydrogen atom covalently bonded to a small, highly electronegativity atoms with lone electron pairs (N, O, F)

hydrogen bond, stronger than dipole dipole (higher boiling point)

Boiling point increases with size apart from hydrogen bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

IMF: London dispersion forces

A

Between nonpolar molecules/atoms

Instantaneous dipole in one particle induces a dipole in another, resulting in an attraction

Larger particles, greater molar mass, more easily polarizable, stronger dispersion forces

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Molecular orbital theory (MO):
When individual atoms combine, atomic orbitals in bonding atoms combine to form molecular orbitals (spread out over entire molecule)
26
Valence bond theory vs. molecular orbital theory
Electrons are localized in VB but in MO they are delocalized over entire atom
27
bond order
+ bond order: # bonding > # antibonding → molecule stable (exists) – or 0 bond order: # bonding # antibonding → molecule unstable (doesn’t exist) Higher bond order = stronger bond
28
Relationship between # electrons in bonding, antibonding molecular orbital & stability
Higher number of electrons in bonding MO = **greater stability** Higher number of electrons in antibonding orbitals = lower stability
29
Percent composition:
% by mass of each element in compound - Elemental makeup of a compound defines its chemical identity - Useful to determine relative abundance of a given element in different compounds of known formula
30
Chemical formula:
relative numbers, not masses, of atoms in substance
31
Empirical formula
derived from experimentally measured element masses by:
32
Empirical formula mass:
avg atomic masses of all atoms in empirical formula
33
Stoichiometry:
quantitative relationships between amount of reactants & products in chemical equation - atoms are neither created nor destroyed - Identity of reactants & products are experimentally determined → cannot change identity (chemical formula) in balancing - Start with more complicated molecules first (greatest #atoms)
34
Limiting reactant:
completely consumed first, limits amount of product formed
35
Theoretical, actual, percent yield
**Theoretical yield**: mass product based on stoichiometry (max) **Actual yield**: mass obtained in the lab **Percent yield** = (actual yield / theoretical yield) * 100
36
Solutions, solute, solvent
**Solutions**: homogeneous mixture of 2+ substances **Solute**: smaller amount **Solvent**: larger amount (ie. water is universal solvent)
37
Why is water an excellent solvent?
Polar nature & ability to form hydrogen bonds Water can dissolve polar & ionic substances
38
Ionic compounds in water:
Forces of attraction between solute & solvent particles dissolve Not all ionic substances are soluble in water Solubility depends on attraction between ions & water
39
Polar covalent compounds: strong, weak, non ELECTROLYTES
Intermolecular hydrogen bonding **Strong electrolytes**: dissociate completely in aqueous solutions (conduct electricity) **Weak electrolytes**: produce fewer number of ions (don’t conduct well) **Non electrolytes**: dissolve in water but don’t produce ions (don’t conduct)
40
Concentration solution:
amount of solute in a given quantity of
41
Molarity:
moles of solute in 1L of solution n = mass / molar mass
42
Dilution: adding additional solvent to a solution
**Dilution**: adding additional solvent to a solution **Dilute**: small amount of solute dissolved **Concentrated**: large amount of solute dissolved
43
Dillution equation
M: molarity V: volume (initial & final)
44
Aqueous chemical reaction types:
45
Aqueous chemical reaction types: Precipitation reaction
**2 soluble ionic compounds** react to give an insoluble product (precipitate) Soluble salts (aq) form clear solutions, insoluble salts form precipitates (s)
46
Aqueous ionic reaction can be expressed in 3 different ways: 1. Molecular equations 2. Ionic equation 3. Net ionic equation
**Molecular equations**: all reactants/products as if they were intact, undissociated compounds Complete **ionic equations**: all soluble ionic substances dissociated into ions **Net ionic equation**: eliminates spectator ions, shows only actual chemical change
47
What are spectator ions?
**Spectator ions**: not involved in actual chemical change (appear unchanged on both sides of ionic equation)
48
Solubility rules for Ionic Compounds in Water
Don’t need to memorize rules, but must know how to interpret
49
Arrhenius definition of acid & base
limited to aqueous solutions (water as solvent) **Arrhenius acid**: acid is a substance that produces H+ ions (protons) in water - hydronium ion = proton = H+ ion **Arrhenius base**: substance that produces OH- ions in water
50
Bronsted-lowry definition of acid & base
much broader definition of acid & base **Acid**: proton donor **Base**: proton acceptor → have a lone pair
50
Can acids and bases act as electrolytes (conduct electricity)?
Strong acids & bases dissociate completely in aq solution → form strong electrolytes Weak acids & bases: dissociate very weakly into ions → form weak electrolytes Reversible reaction (⇌)
50
Neutralization reaction:
acid reacts with a base to form salt & water Complete ionic equation: - strong acids (SA) & strong bases (SB) dissociated - weak acids (WA) & weak bases (WB) (undissociated)
51
Titration: titrant analyte equivalence point
analytical technique to determine concentration of unknown acid/base solution **Titrant**: solution known concentration **Analyte**: solution unknown concentration E**quivalence point**: amount titrant added is enough to completely neutralize analyte solution
52
How to determine the equivalence point?
1. Add **indicator**: change color at/near equivalence point (ie. phenolphthalein) - Color changes at the end point. Ideally, we want the end point to be very close to the equivalent point. 2. Titration curve
53
Relative energies of p-molecular orbitals in homonuclear diatomic molecules (2nd period)
54
Paramagnetic molecules
have at least one unpaired electron
55
diamagnetic molecule
All electrons are paired
56
57
Molecular orbital diagram vs lewis structure in explaining paramagnetic nature of O2 molecule
58
Atomic orbital vs hybrid orbital vs molecular orbital
59
large size and planarity of the naphthalene molecule allow its London dispersion forces to be substantial enough to make it a solid at room temperature, despite these forces being weaker than the hydrogen bonds in water.
60
Bonding orbital
increased electron density directly between nuclei
61
Antibonding orbital
electrons are located away from region between two nuclei (nodal plane)
62
Delocalized pi bond
Pi orbital extend over 2+ atoms In lewis structure, this occurs when with resonance structures involving double & triple bond
63
Molecular orbital diagram (MO)
64
What does a +, 0, negative bond order mean?
+ BO: stable enough to exits ≤ 0: doesn't exit
65
Precipitation reaction
2 soluble ionic compounds (aq) react to give an insoluble product (precipitate (s))
66
Redox reaction
transfer of electron(s) from one species to another even though a solid could be produced, it is not precipitation because the reactants are not soluble salts
67
Oxidation number
assume all bonds are ionic Break bond any two atoms, give both electrons to most electronegative atom (electron transferred from least → most electronegative) Calculate similar to FC (periodic table - atom electrons)
68
Oxidation number rules
69
Identify a redox reaction
calculate oxidation # (state) of each atom before & after reaction
70
OIL RIG
Oxidation: increase in ON (loss electrons) → OIL RIG Oxidizing agent/oxidant Reduction: decrease in ON (gain electrons) Reducing agent/reductant
71
Cation displacement reaction:
metal with high reactivity in the activity series is added to a solution containing a cation with lower reactivity 1. Identify the Reacting Metals: 2. Locate the Metals on the Activity Chart 3. Compare Reactivities: The metal that's higher on the list is more reactive. 3. Predict the Reaction: If the free metal (the one not in a compound) is higher on the activity series than the metal in the compound, the reaction is likely to occur. The free metal will replace the metal in the compound, forming a new compound and releasing the less reactive metal.
72
Atmospheric pressure:
pressure exerted by column if air from top of atmosphere to surface of earth (higher altitude = smaller atmospheric pressure) Barometer device to measure atmospheric pressure → pressure indicated by height (mm) of mercury column [1atm=760mmHg]
73
Gas pressure
74
Boyle's Law
75
Charles law
76
Avogadro's law
77
Ideal gas law
PV = nRT Temperature must be in K
78
Standard temp & pressure (STP):
0ºC, 1atm
79
Standard molar volume:
1 mole of any ideal gas at standard temp & pressure occupies 22.4L
80
Determine molar mass of gas using ideal gas law:
81
Real gasses behave as ideal gasses
high temp and low pressure less influenced by IMF, negligible volume
82
Standard solution: Analyte:
Standard solution: solution known concentration Analyte: solution unkown concentration
83
Equivalence point (indicator) in acid, base titration
moles H+ from acid = moles OH- from base amount H+ ions in flask = amount OH- ion added
84
Endpoint point (indicator) in acid, base titration
slight XS of base (OH-), indicator changes color
85
Diprotic acid
donate two hydrogen ions (H⁺) per molecule in an aqueous solution

Tufts - Chem 1 (4 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2025 Bold Learning Solutions. Terms and Conditions