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Final exam review Flashcards

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1
Q

What is a covalent bond?

A

A sharing of electrons between two non-metal nuclei

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2
Q

What is an ionic bond?

A

A transfer of electrons between a metal and a non-metal atom

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3
Q

What is metallic bonding?

A

Why most metals are solids

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4
Q

What do lewis dot structures represent?

A
  • Valence electrons around an atom
  • H follows the duet rule, it forms stable
    compounds by sharing 2 electrons
  • Most main-group atoms follow the octet rule:
    they tend to have 8 electrons (bonds + lone
    pairs) in their valence shell
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5
Q

How do you quantify a bonds strength?

A

Ionic bonds - lattice energy
Covalent bonds - bond energies

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6
Q

What is lattice energy?

A

The delta H when gas phase ions form one mole of an ionic solid
- Always exothermic
- More exothermic lattice = stronger ionic bonds
- Calculated using Hess’ Law
- 3A + B = C, C = 2D so 3A + B = 2D
- H1 H2 (H1 + H2)

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7
Q

What does it mean to use the Born-Haber Cycle?

A

Using Hess’ Law to calculate lattice energies (weaker ionic bonds)
- k(Q1Q2/r)
- More exothermic LE’s have the elements with
the larger charge (BaTe over LiF), the larger
element (FeCl3 over FeCl2), or are smaller (ScCl3
over ScBr3)

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8
Q

How do you show polar covalent bonds?

A
  • Partial charges
  • Dipole; arrow points to more electronegative
    atoms
  • Electrostatic potential map, more red means the
    region is more electron-rich and more blue
    means the region is more electron-poor
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9
Q

What is electronegativity?

A

The ability of an atom to attract shared electrons to itself in a chemical bond
- If delta EN is small (< 0.4) the bond is covalent
(example: Cl2)
- If delta EN is medium (0.4 - 2.0) the bond is polar
covalent (example: HCl)
- If delta EN is large (> 2.0) the bond is ionic
(example: NaCl)

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10
Q

What are the characteristics of an ionic compound?

A
  • Large delta EN
  • Cation-anion arrays
  • Solids at STP
  • High melting/boiling points
  • Conduct as liquids
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11
Q

What are the characteristics of a covalent compound?

A
  • Small delta EN
  • Discrete molecules
  • Solid/liquid/gas at STP
  • Low melting/boiling points
  • Insulating as liquids
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12
Q

What is a dipole moment?

A
  • A measure of positive and negative charge
    separation in a molecule
  • Dipole moment is the vector sum of individual
    bond dipoles
    • EN is calculated; DM is measured by the
      current flow when the sample is between the
      charged plates of a capacitor
    • Molecules have a DM if they can be oriented
      such that one side has a positive partial
      charge and the other has a negative one
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13
Q

How do you draw a lewis structure?

A
  • Count valence electrons from all atoms
  • Connect atoms using single bonds
  • Leftover electrons go to multiple bonds and lone
    pairs to follow duet (H) and octet rules
  • The central atom is usually the least
    electronegative
    • Exception: H is never the central atom
  • If there is one of a certain type of atom, it is
    usually the central atom
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14
Q

What does it mean for a molecule to have resonance?

A
  • > or = 2 valid Lewis structures for one species
  • The real structure is an average of the resonance
    forms
    • Resonance structures aren’t in equilibrium
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15
Q

What are the two types of charge?

A
  • Ionic: on entire ion
  • Formal: on each atom
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16
Q

What is a formal charge?

A
  • Oxidation state completely assigns electrons in a
    bond to the more electronegative element
  • Formal charge divides electrons equally, ignoring
    electronegativity
    • Formal charge = (valence electrons on neutral
      atom) - [non-bonding electrons + (1/2 *
      bonding electrons)]
    • There are 2 non-bonding electrons in 1 lone
      pair
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17
Q

How do you solve for formal charge?

A
  • For a molecule: E(FCs) = 0
  • For a polyatomic ion: E(FCs) = its charge
  • FCs allow us to evaluate resonance structures:
    • Good resonance structures have low units of
      FC (not ) is best)
    • Negative FCs should be on the most
      electronegative atom; positive FCs should be
      on the least EN atom
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18
Q

What are the octet rule exceptions?

A
  • Second-row elements can’t exceed an octet
    since 2d orbitals don’t exist
  • Third-row and heavier elements often obey the
    octet rule, but can use empty d orbitals to
    exceed it
  • Some main group atoms typically hold fewer
    than 8 electrons, (example Be (4), B (6) and Al (6))
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19
Q

What do you prioritize when finding the best resonance form of a structure?

A

Fewer formal charge units > obeys octet rule

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20
Q

What is a bond energy?

A

The energy required to break the bond of a gas-phase molecule
- Endothermic; it always takes energy to break a
bond
- Compounds with stronger bonds are more
stable and less chemically reactive
- General bond strength trends: X3Y > X=Y > X-Y
- BE’s are easy to obtain for molecules with only
one type of bond

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21
Q

What is average bond energies?

A

An average of the measured bond energies for a particular bond in several compounds

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22
Q

What can delta H rxn be estimated as?

A

Delta H rnx = E B.E. (bonds broken) - E B.E. (bonds formed)
- B.E. (bonds broken) - energy consumed
- B.E. (bonds formed) - energy released

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23
Q

What are the average bond lengths?

A
  • Triple (shortest) < double < single (longest)
  • Length decreases left to right across a period
  • Length increases down a group
  • In general, as bonds get longer, they also get
    weaker
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24
Q

What is VSEPR theory?

A

Valence Shell Electron Pair Repulsion
- Predicts the shape of covalent compounds
- Regions of electron density orient around a
central atom as far apart as possible

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25
Q

What is molecular geometry?

A
  • The shape of covalent compounds is described
    by regular polygons and Platonic solids
    - Trigonal planar
    - Square planar
    - Tetrahedral
    - Octahedral
  • The central atom is placed in the center
  • The corners of the shape are the positions of
    surrounding atoms
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26
Q

What are electron groups?

A
  • Molecular geometry is found by counting
    electron groups on the central atom
  • These count as one electron group:
    - Lone pairs
    - Single bonds
    - Multiple bonds
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27
Q

What is the corresponding electron geometry for each electron group?

A
  • 2 groups = linear geometry
  • 3 groups = trigonal planar geometry
  • 4 groups = tetrahedral geometry
  • 5 groups = trigonal bipyramidal geometry
  • 6 groups = octahedral geometry
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28
Q

Characteristics of two electron groups

A
  • Linear
  • A central atom with two electron groups has
    linear electron geometry
  • Bond angle is 180
    - Examples: NO2, BeF2, HCN, SCN, CO
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29
Q

Characteristics of three electron groups

A
  • Trigonal planar
  • A central atom orients three electron groups 120
    apart and coplanar
    - Examples: SO2, SO3, NO3, NO2, BF3
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30
Q

Characteristics of four electron groups

A
  • Tetrahedral
  • A central atom puts four electron groups 109.5
    apart in a tetrahedral orientation
    - Examples: PO4, NH4, CF4, ClO4
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31
Q

Characteristics of five electron groups

A
  • Trigonal bipyramidal
  • A central atom with five electron groups adopts
    a trigonal bipyramidal orientation
  • 90, 120, and 180 bond angles
    - Examples: PF5, SF4, I3, ClF3
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32
Q

Characteristics of six electron groups

A
  • Octahedral
  • A central atom with six electron groups has
    octahedral electron geometry
  • 90 and 180 bond angles
  • All positions are equivalent
    - Examples: PF6, SF6, XeCl4, BrF5
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33
Q

What are lone pairs?

A
  • Lone pairs have hemispherical electron density
    regions
  • Bonded electrons (not lone pairs) have
    cylindrical electron density regions
  • Lone pairs occupy more space around an atom
    than bonded electrons
  • Since lone pairs take up more space than
    bonding electrons, they force atoms closer
    together
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34
Q

What is AXE notation?

A
  • A = central atom
  • X = outer atoms
  • E = lone pairs on central atom
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35
Q

AX2 and AXE

A

Two electron groups on a central atom, whatever they are, is always linear molecular geometry

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36
Q

Ax2E

A
  • Electron geometry = trigonal planar
  • Molecular geometry = bent
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37
Q

AX3E

A
  • Electron geometry = tetrahedral
  • Molecular geometry = trigonal pyramidal
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38
Q

AX2E2

A
  • Electron geometry = tetrahedral
  • Molecular geometry = bent
39
Q

AX4E

A
  • Electron geometry = trigonal bipyramidal
  • Molecular geometry = see-saw
40
Q

AX3E2

A
  • Electron geometry = trigonal bipyramidal
  • Molecular geometry = T-shaped
41
Q

AX2E3

A
  • Electron geometry = trigonal bipyramidal
  • Molecular geometry = linear
42
Q

AX5E

A
  • Electron geometry = octahedral
  • Molecular geometry = square pyramidal
43
Q

AX4E2

A
  • Electron geometry = octahedral
  • Molecular geometry = square planar
44
Q

What are the two main bonding theories?

A
  • Valence bond theory
    - Atomic orbitals mix to become hybrid
    orbitals between two atoms
  • Molecular orbital theory
    - Atomic orbitals mix to become molecular
    orbitals that span the entire molecule
45
Q

What is hybridization?

A

Mixing atomic orbitals (s, p, d, …) to form a new set of degenerate hybrid (sp, sp2, …) orbitals
- x atomic orbitals mix to x hybrid orbitals
- Hybridization explains the observed shapes of
molecules

46
Q

Overview of hybridization

A
  • 2 groups, sp hybridization, linear geometry
  • 3 groups, sp2 hybridization, trigonal planar
    geometry
  • 4 groups, sp3 hybridization, tetrahedral
    geometry
  • 5 groups, dsp3 hybridization, trigonal
    bipyramidal geometry
  • 6 groups, d2sp3 hybridization, octahedral
    geometry
47
Q

sp3 hybridization

A
  • One s and three p orbitals to give four sp3
    orbitals
  • sp3 orbitals have equivalent shapes and
    energies
  • Each points to a different vertex of an imaginary
    tetrahedron
48
Q

sp3 orbitals

A
  • sp3 orbitals are higher in energy than their
    parent s orbitals but lower in energy than p
    orbitals
  • Unlike p orbitals, hybrid orbitals have lobes of
    different sizes
  • The larger lobe overlaps with other orbitals for
    bonding
49
Q

sp2 hybridization

A
  • For three electron groups, one s and two p
    orbitals hybridize to three sp2 orbitals
  • Trigonal planar geometry is obtained by
    hybridization of one s and two p orbitals to give
    three sp2 orbitals
  • One 2p orbital remains unhybridized
50
Q

sp2 orbitals

A
  • Trigonal planar geometry is obtained by
    hybridization of one s and two p orbitals to give
    three sp2 orbitals
  • One 2p orbitals remains unhybridized
51
Q

What are sigma bonds?

A
  • Head-to-head orbital overlap
  • Electron density between nuclei
  • All orbitals form sigma’s
52
Q

What are pi bonds?

A
  • Side-to-side orbital overlap
  • p orbitals on adjacent atoms overlap
  • Electron density above and below the
    internuclear axis
53
Q

What types of bonds are contained within single, double and triple bonds?

A
  • Single bond = one sigma bond
  • Double bond = one sigma and one pi bond
  • Triple bond = one sigma and two pi bonds
54
Q

sp hybridization

A
  • sp hybridization is for two electron groups on a
    central atom
  • One s and one p orbital hybridize to two sp
    orbitals
  • Two p orbitals are unchanged
55
Q

What are the problems with Lewis Theory?

A
  • Often requires > or = 2 structures to describe
    bonding (resonance
  • Doesn’t predict bond strength
  • Doesn’t predict bond length
  • Doesn’t predict magnetic behavior
  • Doesn’t describe odd-electron species well
56
Q

What can be done with molecular orbitals?

A
  • The wavefunctions of two atomic orbitals can:
    - be added, forming a bonding MO
    - be subtracted, forming an antibonding MO
  • x atomic orbitals produce x MOs
57
Q

Bonding vs. Antibonding MOs

A
  • Electrons in bonding MOs are stabilizing
    • Lower energy than atomic orbitals (AOs)
    • Electron density is between nuclei
  • Electrons in antibonding MOs are destabilizing
    • Symbolized by *
    • Higher energy than AOs
    • Electron density is outside the internuclear
      axis
    • Electrons in antibonding orbitals cancel
      stability gained by electrons in bonding
      orbitals
58
Q

How do you use MO diagrams?

A
  • Select correct diagram
  • Count the number of valence electrons
  • Add valence electrons
    - Fill lowest energy MOs first (Aufbau)
    - Each MO holds 2 electrons (Pauli Principle)
    - Half-fill degenerate orbitals (Hund’s rule)
59
Q

How do you find bonding order (BO)?

A

bond order = (bonding electrons - antibonding
electrons) / 2
- BO = 0: the bond is unstable and will not form
- BO > 0: the bond will form
- The larger the BO, the stronger and shorter the
bond

60
Q

What are the two types of p orbital overlap?

A
  • Head to head in a sigma bond
  • Side to side in a pi bond
61
Q

What are the two types of MO overlaps?

A
  • Bonding - between nuclei
  • Antibonding - outside nuclei
62
Q

What is photoelectron spectroscopy?

A
  • A molecule is bombarded with X-rays to remove
    an electron
  • The electron’s energy is measured, giving the
    relative energy levels
63
Q

Inter vs Intramolecular forces

A
  • Intermolecular - attractions between > or = 2
    species
    - Hydrogen bonds (4-50 kJ/mol of strength)
    - Dipole-dipole (5-20 kJ/mol of strength)
    - London force (0.1-5 kJ/mol of strength)
  • Intramolecular - hold atoms together in a
    compound (covalent or ionic bond)
    - Ionic bond (400-500 kJ/mol of strength)
    - Covalent bond (100-400 kJ/mol of strength)
64
Q

What are intermolecular forces (IMFs)?

A
  • IMF strength determines phase:
    • Solids have lots of strong IMFs
    • Liquids have fewer and weaker IMFs
    • Gases have none (theoretically)
65
Q

What are dipole dipole forces?

A
  • Electrostatic attractions between partial charges
    of polar molecules
  • One molecule’s positive pole attracts the
    negative pole of another
66
Q

What is hydrogen bonding?

A
  • A strong dipole-dipole force that occurs in
    molecules with H-X bonds
  • X can only be N, O and F because they are:
    • Small, molecules approach closely
    • Electronegative, strong partial charges on H
      and X
67
Q

What are dispersion forces (london)?

A
  • Present in all atoms, molecules and ions
  • Responsible for liquid and solid phases of
    nonpolar molecules
  • A temporary charge build-up at one side induces
    a dipole in its neighbor
68
Q

What is polarizability?

A
  • How easily a species’ electron cloud distorts
  • More polarizable species:
    - From stronger dispersion forces
    - Have higher molar masses
    - Have greater surface areas
  • In general: the stronger the IMFs, the higher the
    boiling point (bp)
69
Q

What is boiling point?

A

The temperature where P vapor = P atmosphere
- When a liquid reaches its bp, its particles have
enough energy to break IMFs
- To boil a liquid, energy is needed to break all
IMFs between the particles
- Higher boiling points result from:
- Strong IMFs
- Large molar mass
- Shape

70
Q

What is ion-dipole forces?

A

An IMF between an ion and the oppositely charged end of a polar molecule

71
Q

What is vaporization?

A
  • Vaporization or evaporation - liquid-to-gas phase
    change below the liquids boiling point
  • Endothermic; energy is required to break a
    liquid’s IMFs
  • Heat of vaporization (delta H vap) - energy
    required to vaporize 1 mol of liquid at 1 atm
72
Q

What is dynamic equilibrium?

A
  • In a closed container a liquid’s volume decreases
    then remains constant
  • At equilibrium: evaporation rate = condensation
    rate
73
Q

What is vapor pressure?

A

The partial pressure of vapor over a liquid or solid at equilibrium

74
Q

What is a vapor?

A

The gas phase of a substance that is solid or liquid at 25 C and 1 atm

75
Q

How does vapor pressure increase with temperature?

A
  • P vap increases exponentially with temperature
  • At higher T, more particles have the energy to
    enter the gas phase, so P vap increases
76
Q

What is heat of fusion?

A

The energy required to melt one mole of a solid

77
Q

What is heat of sublimation?

A

The energy required to sublime one mole of a solid

78
Q

What is the heating curve of water?

A
  • T is measured as a sample is heated
  • T is constant during phase changes
  • Slope differ; Cs (ice) not equal to Cs (water) not
    equal to Cs (steam)
79
Q

What are phase diagrams?

A
  • Phase diagrams display phases as a function of
    temperature and pressure
  • They are unique to each substance
  • Most substances have at least three phases
80
Q

What is the triple point?

A
  • At the triple point all three phases are in
    equilibrium
  • The triple point of water occurs at 0.01 C and
    0.0060373 atm
  • It can only be observed in sealed containers
    under a high vacuum
81
Q

What is a phase diagram?

A
  • Ask first if the substance is denser as a liquid or
    a solid
  • As a solid; its liquid - solid boundary has a
    positive slope
82
Q

What are supercritical fluids?

A
  • SFs fill their container and diffuse through solids
    (like a gas) and dissolve solids (like a liquid)
  • Beyond the critical point, the meniscus
    disappears
  • Supercritical CO2 is used to remove caffeine
    from coffee, nicotine from tobacco, fat from raw
    meat, dirt from clothes (dry cleaning), and
    pesticides from fruit
83
Q

What kind of solvent will a solute be dissolved in?

A
  • A solute will dissolve in a solvent if both have
    similar IMF types
  • Solvent molecules will attract the solute particles
    at least as well as the solute particles to each
    other
84
Q

What can affect solubility?

A
  • Even molecules with hydrogen bonding become
    water-insoluble if the rest of the molecule only
    has London forces
  • The solubility of most salts increases with
    temperature
  • Solubility: the amount of a substance that will
    dissolve in a given amount of solvent
85
Q

Solvent vs solute

A
  • Solvent: the majority component of the solution
  • Solute: the minority component of the solution
    • Example: mouthwash
      - Solvent = water
      - Solute = ethanol
    • Example: absinthe
      - Solvent = ethanol
      - Solute = water
86
Q

Molality vs Molarity

A
  • Both use moles of solute
  • Molarity uses solution volume, molality uses
    solvent mass
  • Molarity is temperature-dependent, molality is
    not
  • They are close for dilute aqueous solutions
    • molarity (M) = mol solute / L solution
    • molality (m) = mol solute / kg solvent
87
Q

What is colligative properties?

A
  • Solution properties that depend on the number
    of solute particles, not their identity
  • The colligative properties are:
    • Vapor pressure lowering
    • Boiling-point elevation
    • Freezing-point depression
    • Osmotic pressure
88
Q

What is boiling point elevation?

A
  • A nonvolatile solute raises a solvents boiling
    point
  • The solute can be an electrolyte or a
    nonelectrolyte
    - Delta T b = boiling point elevation
    - K b = bp elevation constant
    - m solute = solute molality
    - T = K * m
89
Q

What is freezing point depression?

A
  • Any nonvolatile solute raises a solvent’s boiling
    point and lowers its freezing point
  • The solute can be an electrolyte or a
    nonelectrolyte
  • For a given solvent: K f > K b
90
Q

What is osmosis?

A

Solvent flow through a membrane from lower to higher concentration

91
Q

What is osmotic pressure?

A

The pressure that must be applied to a solution to stop osmosis

92
Q

What is the van’t Hoff Factor?

A
  • The ratio of moles of particles in solution to
    moles of solute
  • For nonelectrolyte, the van’t Hoff factor is 1
  • For an electrolyte, the van’t Hoff factor is the
    number of ions formed in solution
93
Q

What is an ion pair?

A
  • An IMF between a dissolved anion and cation
  • An ion-pair acts as a single solute particle
  • Ion pairing affects colligative properties by
    reducing the number of solute particles
  • For electrolytes, observed i < expected i
  • All electrolytes ion-pair, but less so in dilute solutions

Chemistry (9 decks)

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