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Year 1 Finals > Chemistry 101 Final Unit 3 > Flashcards

Chemistry 101 Final Unit 3 Flashcards

1
Q

What are gases’ physical properties?

A
  • largely independent of their chemical composition
  • highly compressible (Boyle’s law), thermally expandable (Charles’ law), low viscosity (resistance to deformation), low density, infinitely miscible (a mixture of gases)
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2
Q

What is the ideal gas law formula?

A

PV: nRT
- P: pressure (atm) (1 atm: 760 torr: 760 mmHg)
- V: volume (L)
- T: absolute temperature (K)
- n: amount of gas (mol)
- R: gas constant (0.08206 atm/molxK)

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3
Q

What are the 3 applications of the ideal gas equation?

A
  1. molar mass determination
  2. gas density
  3. gas stoichiometry
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4
Q

What is the kinetic molecular theory?

A

Model of ideal gas (explains behaviour of gas)
1. particles are in random motion
- collisions of particles with the walls are the cause for
pressure exerted by the gas
2. negligible particle volume
- volume of individual particles can be neglected (rounds to
0), as gas molecules are small compared to the distance
between them
3. particles collide with each other + container wall
4. particles collide with each other + experience no interparticle forces
- no attractive or repulsive forces between particles
5. constant total energy
- but energy is transferred in collisions with KE conserved

average KE of collection of gas particles: gas’ temperature

average KE of an ideal gas only depends on its temperature

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5
Q

What are the 2 relationships between gases and temperatures?

A
  1. different gas + same temperature
    • lighter gases: greater average speeds (look at atomic mass)
  2. same gas, different temperature
    • higher temperature: greater average speeds (increase in KE:
      increase in average speed)
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6
Q

What is the relationship between temperature, KE and motion?

A

Higher temperature: greater KE: greater motion
- note that lighter mole will have greater molecular speed

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7
Q

What is ideal and non-ideal for gases?

A
  • ideal: lower pressure, higher temperature
  • non-ideal: higher pressure, lower temperature
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8
Q

What occurs in real gases that violate the assumptions made in KMT?
1. no stickness: no interparticle forces
2. no size: no negligible particle volume

A
  1. particles in a real gas experience weak interparticle attractions
    • interparticle attractions occur between separate atoms or
      molecules + are caused by imbalances electron distribution,
      these forces are important only over very short distances +
      are much weaker than bonding forces
  2. particles in a real gas occupy a finite volume
    • at normal Pext, the space between particles is enormous
      compared with the volume of the particles themselves
    • increase Pext: decrease in free volume thus the volume of
      particles make up a significant portion of the container
      volume
    • at moderately high Pext, z values are lower than ideal
      values (2>1) due to interparticle attractions at very high
      Pext, values are greater than ideal (2>1) due to particle volume
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9
Q

What is van der Waals equation used for?

A

Used to account for the non-ideal behaviour.

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10
Q

What are phase changes?

A

When real gas’ behaviour deviates from being ideal when attractive forces between gas particles become significant, when these forces are strong enough: ex. gas becomes liquid.
- LDF: hold particles together, weaker than covalent bonds
- phase: a physically distinct homogeneous part of system

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11
Q

What are intramolecular forces + intermolecular forces?

A
  • intramolecular (bonding) forces
    • within a molecule, chemical behaviour of 3 states are all
      identical since they consist of the same molecule held
      together by the same covalent bonding forces
  • intermolecular (antibonding) forces
    • between molecules, physical behaviour of the states differ
      since strength of forces differ
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12
Q

What do phase changes depend on?

A
  • intermolecular forces + KE of the moving particles
    • T increase: KE average increase (particles move faster +
      overcome attractions more easily)
    • T decrease: KE average decrease (particles move slower +
      attractions can pull them together closer more easily)
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13
Q

Substance’s stage changes can absorb or release energy, what are the 6 changes?

A
  1. s - l - g: energy’s absorbed (△H° fus, △H° vap)
    • s-l: melting, l-g: vaporization, s-g: sublimation
  2. g - l - s: energy’s released (-△H° fus, -△H° vap)
    • g-l: condensation, l-s: freezing, g-s: deposition
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14
Q

Which one is greater, △H° fus, △H° vap?

A

△H° fus «< △H° vap in general for all substances
- less energy to overcome the IMFs enough for molecules to
move out of fixed positions (melt a solid) than to fully
separate them (vaporize a liquid)

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15
Q

Why is H2O a special case for phase changes?

A

Both △H° fus + △H° vap are large: stronger IMFs are needed to be overcome + hydrogen bonding.

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16
Q

What is the equilibrium process?

A

Eventually rate of vaporization equals rate of condensation: equilibrium’s established.
- molecular level: molecules enter + leave at equal rates

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17
Q

What is vapour pressure?

A

The measure of the tendency of a substance to change into the gaseous/vapour state from the liquid state.

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18
Q

What is temperature’s effect on vapour pressure?

A
  • changes the fraction of molecules moving fast enough to escape liquid: fraction slow to be recaptured
  • T increase: vp increase: T increase leads to more molecules having enough energy to leave the surface
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19
Q

What does vapour pressure depend on?

A
  • type of liquid + temperature
    • molecules with weaker IMF forces are held less tightly +
      vaporize more easily
    • molecules with stronger IMF forces are held more tightly +
      vaporize less easily
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20
Q

What is the relationship between vapour pressure + boiling point?

A

Substances with vp increase: bp decrease, which leads to volatile (tendency of substance to evaporate at normal temperature).

21
Q

What is the boiling point?

A
  • temperature at which the vapour pressure (inside bubbles) equals the external pressure (usually atmospheric pressure).
  • normal bp: temperature where vp: 1 atm
22
Q

For solid-liquid equilibria…

A
  1. T increase: particles vibrate more rapidly until some have enough KE to break free of their positions: melting begins
  2. more liquid molecules, some collide with solid and become fixed in position again
    • phases remain in contact: dynamic equilibrium when
      melting rate: freezing rate - the T when this happens is mp
  3. liquids + solids are nearly incompressible since pressure has small effect on the melting + freezing rates
23
Q

For solid-gas equilibria…

A
  1. substance sublimes, not melts since intermolecular forces aren’t great enough to keep molecules close when they leave solid state
  2. some solids do have high enough vp to sublime at ordinary conditions (ex. dry ice is non-polar: weak IMF)
24
Q

What are phase diagrams?

A
  • depicts which phase is most stable (at different temperature + pressure conditions)
  • regions: indicates which one phase is most stable
    • generally, solids are stable at low T + high P while gases are
      stable at high T + low P
  • lines: indicates which two phases coexist in equilibrium
    • generally, s is denser than l: increase in converts, l to s
      (water’s the exception)
  • critical point: where liquid-solid line terminates
    • above Pc + Tc, the vapour can’t be liquified: supercritical
      fluid
    • when a liquid is heated in a closed container, its density
      decrease and at the same time more of the liquid vaporizes
      so the density of the vapour increases
  • at cp: the densities become equal + phase boundary
    disappears, KE average is too high that vapour can’t be
    condensed at any pressure
  • triple point: when solid, liquid + gas can coexist in equilibrium
25
Q

What are intermolecular forces?

A
  1. only influences physical (not chemical) properties of a substance
  2. strengths are in: mp + bp, heats of fusion + vaporization, surface tension + viscocity, solubility
  3. not all substances are molecular
    • no molecules in simple ionic solids, extended network
      solids
  4. van der Waal forces
    • nonbonding forces, very weak compared to chemical
      bonding forces
26
Q

What are the different types of chemical forces?

A
  1. bonding
    • ionic: cation - anion
      • ions are attracted to other ions with opposite charge
    • covalent: nuclei + shared e-pair
    • metallic: cations - delocalized electrons
  2. non-bonding or van der Waals
    • ion-dipole: ion charge - dipole charge
    • dipole-dipole: dipole charges
      • molecules with permanent dipole moments experience
        dipole-dipole interactions
      • molecules will orient to maximine + & - attractions +
        minimize like-charge repulsion
    • H-bond: polar bond to H-dipole charge
      • special type of dipole-dipole between molecules that have
        H atom bonded to small, very EN atom (N, O or F)
    • ion-induced dipole: ion charge - polarizable e-cloud
    • dipole-induced dipole: dipole charge - polarizable e-cloud
    • london dispersion (induced dipole - induced dipole):
      polarizable e-clouds
      • caused by molecule’s polarizability, which temporarily
        makes that point of molecule negatively charged with the
        rest of the molecule positively charged
27
Q

What is polarizability?

A
  • how easy it is to cause distortion to the e-cloud to create a dipole moment
    • nonpolar molecule: the distortion induces a dipole moment
      polar molecule: the distortion induces an increase in the
      existing dipole moment
    • increase # of electrons + increase molecular size: increase
      polarizability (boys + girls example)
      • increase down the group: since atomic size increases +
        larger e-clouds easier easier to distort
      • decrease across a period: increase Zeff: moves the atoms
        smaller + holds electrons more tightly
      • cations are less polarizable (smaller, less electrons) while
        anions are more polarizable (bigger, more electrons)
28
Q

What are liquids?

A
  • have definite volume but aren’t confined to shape of container, properties depend on IMF’s strengths
29
Q

What is surface tension?

A
  • liquid’s resistance to increase its surface tension, the energy required to crease the surface tension by a given amount
  • depends on IMF present
  • when considering the forces acting on a particle at the surface VS in the interior
    • interior molecule: attracted by others on all sides
    • surface molecule: attracted by others below + sides, so it
      experiences a net attraction downward
30
Q

How do you get a high surface tension?

A

Strong IMF: the stronger the IMF between particles, the more energy it takes to increase SA, so the greater the SA

31
Q

What is the relationship between surface tension and temperature?

A

SA decrease: T increase (at higher T, the liquid molecules have increased average KE with which to break attractions to molecules in the interior)

32
Q

What is capillarity and wetting?

A
  • capillarity: rising of a liquid in a tube
  • wetting: spreading of a liquid across a surface
33
Q

Consider competition between cohesive forces (within the liquid) and adhesive forces (between liquid + surface)…

A
  • cohesive > adhesive: retains shape
  • cohesive < adhesive: liquid spreads out
34
Q

Ex. water + mercury…-

A
  • water wets glass: mensicus is concave (molecules of the liquid are attracted to those in the container)
    • adhesive forces (glass-H2O) are stronger than cohesive
      forces (H2O-H2O) bottom of meniscus is below the H2O-
      glass contact line
  • mercury doesn’t wet glass: meniscus is convex (molecules of the liquid are attracted to each other)
    • cohesive forces (Hg-Hg) are stronger than adhesive forces
      (Hg-glass) top of mensicus is above the Hg-glass contact line
35
Q

What is viscocity?

A

Resistance of a liquid to flow, strong increase (ex. honey + motor line) implies strong IMF.
Affected by molecular shape.

36
Q

What is the relationship between temperature and viscosity?

A

Increase T: decrease viscosity: faster moving particles overcome IMF more easily, thus resistance to flow decreases with temperature increases.

37
Q

What is solubility?

A
  • ability to be dissolved
  • solution (homogeneous mixture), solute (minor component, is dissolved), solvent (major component, the one doing the dissolving)
  • attractive forces between same molecules, ions (in solid) are replaced by similar attractive forces with solvent molecules (ex. solute-solvent forces)
38
Q

What are solids?

A
  • have definite shape + volume (particles are tightly packed + barely move)
  • can be amorphous (poorly defined shapes due to their particles lack orderly arrangement) or crystallin solids (well defined shapes due to their particles occurring in an orderly arrangement)
39
Q

How do we describe the structure of crystals?

A
  • by specifying a pattern + the atoms that belong to the pattern
    • highly symmetrical external form of crystals suggest regular
      ordered arrangement of atoms, we describe the contents
      of the unit cell, which then translated, reproduces entire
      crystal structure
  • lattice: ordered array of points that describe particles’ arrangement that form a crystal
    • lattice is imaginary, the crystal structure is what’s real
40
Q

What are the 3 different types of cubics?

A
  1. simple cubic: 8 identical particles define the corners
  2. body-centered cubic: 8 identical particles define the corners, 1 at the center
  3. face-centered cubic: 8 identical particles define the corners, 6 at the center of each face
41
Q

What is coordination number?

A

The number of nearest neighbours of a particle in a crystal.

42
Q

How are different parts of a cube shared?

A
  1. corner is shared by 8 cells: 1/8 contribution
  2. face is shared by 2 cells: 1/2 contribution
  3. edge is shared by 4 cells: 1/4 contribution
  4. particles inside a cell are not shared: one whole atom or particle
43
Q

What are the 2 different packings?

A
  1. hexagonal closet packing
    • coordination number: 12
    • 6 atoms per unit cell
  2. cubic closest packing
    • coordination number: 2
    • 4 atoms per unit cell
44
Q

What are the 3 common types of structures found for metallic elements?

A
  1. hcp
    • stacking sequence: AB AB…
    • lattice: primitive hexagonal
  2. fcc
    • stacking sequence: ABC ABC
    • lattice: face-centered cubic
  3. bcc
    • stacking sequence: (not closest-packed)
    • lattice: body-centered cubic
45
Q

What do combinations of 2 or more metals give?

A

Improved physical properties (strength, hardness, resistance to corrosion).

46
Q

What are the 3 ionic solids?

A
  1. sodium chloride structure (rock salt)
    • face-centered cubic array of anions with an interpenetrating
      fcc cation lattice
    • the smaller Na+ ions end up in the holes between the larger
      Cl- ions; thus, each Na+ is surrounded by 6 Cl- & vice versa
  2. zinc blende structure
    • zinc lattice is best thought as fcc arrangement of anions,
      occupying one half of the tetrahedral holes
    • each ion is 4-coordinate and not a tetrahedral geometry
  3. fluorite structure
    • common among salts with a 1/2 cation/anion ratio that
      have relatively large cations + small anions
    • antifluorite structure is common in compounds with a 2/1
      cation/anion ratio and a relatively large anion, the ion
      arrangement is the opposite of that in the fluorite structure
47
Q

Carbon occurs in several elemental forms (allotropes) diamond + graphite. What are the differences?

A
  1. diamond
    • bonding: adopts fcc unit cell with each C tetrahedrally
      bound to 4 others in an endless array, electrons are
      localized
    • properties: transparent, poor thermal + electrical
      conductors
    • uses: jewelry, drilling
  2. graphite
    • bonding: occurs as stacked flat sheets of hexagonal C rings
      with a strong sigma-bond framework + delocalized pi-bonds
    • properties: soft + slippery, good electrical conductivity
    • uses: pencils, lubricants, dry cell
48
Q

What are fullerenes?

A

Group of carbon allotropes based on large molecular clusters of carbon atoms; prototype is the C60 molecule.

49
Q

What are the 3 components of the ideal gas law?

A
  1. Boyle’s Law describes relationship between volume and pressure, states that volume is inversely proportional to temperature.
  2. Charles’ Law describes the relationship between temperature and volume, states that volume is directly proportional to temperature.
  3. Avogadro’s Law states that volume is directly proportional to amount.

Year 1 Finals (4 decks)

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