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Chem 103 > Exam 4 > Flashcards

Exam 4 Flashcards

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1
Q

Recap: What Do We Know?

A
  • When atoms form new compounds, the properties are emergent
    (not the sum of the atoms).
  • Properties of materials depend upon the types of bonds, spatial
    arrangement of atoms, and the interactions between molecules.
  • Atoms interact electrostatically – interactions range from IMFs to
    bonding to ion-ion interactions.
  • The way atoms interact depends on the arrangement of electrons
    (valence electrons).
  • When atoms interact, the system becomes more stable (and
    releases energy to surroundings).
  • All attractive interactions require energy to overcome.
  • Properties (such as mp and bp) allow us to make predictions about
    types of interactions that are present.
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2
Q

properties of Metallic bonding

A

high melting point
-> break metallic bonds
conductors
-> delocalized electrons

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3
Q

properties of a covalent network

A

high melting point
-> break covalent bonds
mostly insulators
-> electrons localized in bonds

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4
Q

properties of discrete molecules

A

lower melting points
- only disrupt IMFs
insulators
-> electrons localized in bonds

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5
Q

properties of NaCl

A
  • forms colorless crystals that are often cubic in shape
  • hard and brittle
  • high melting point
  • conducts electricity when melted (but not as a solid)
    – Can this be explained with any of our current models of bonding? NO
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5
Q

What type of compound is sodium chloride

A

ionic compound (Na+ alternating Cl-)

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6
Q

the continuum of bond types

A

(from lowest electronegativity differences to highest)
pure nonpolar covalent bonds (ie H-H)
polar covalent bonds (Ie Cl-H)
ion-ion interactions (no such thing as an ionic bond) (an electrostatic attraction) (ie cl- na+)

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7
Q

characteristics of simple ionic compunds

A

-contain a metal and a nonmetal (eg NaCl)
-> metal has low electronegativity and nonmetal has high electronegativity
- ionic compounds are neutral
- Metals form cations and nonmetals form anions
-> each ion typically achieves noble gas configuration

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8
Q

why do ionic compounds form?

A
  • requires energy to form ions (either adding or removing electrons)
  • ion-ion interactions release energy and stabilize system
  • ionic compounds from when ion-ion interactions release more energy than is needed to form ions
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9
Q

cations

A

-positively charged
- simple cations form when metals lose electrons
- charge usually corresponds to loss of electrons back to “core”
- ions are less stable than neutral atom- loss of electrons always requires energy
- why do cations form then?
-> because more energy is released then ion-ion interactions are formed

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10
Q

Anions

A

-negatively charged
-simple anions formed when non-metals gain electrons
- charge usually corresponds to gain of electrons to the next filled (sub) shell (ie noble gas configuration)
- ions are less stable than neutral atoms- gaining electrons always requires energy
- why do anions form?
-> because more energy is released when ion-ion interactions are formed

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11
Q

which is bigger; Na atom or Na+ cation

A

Na atom because Na+ has same number of protons as Na but less electrons meaning a greater attractive force between nucleus and electrons and less electron-electron repulsion

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12
Q

which is bigger; Cl- anion or Cl atom

A

cl- anion because Cl- has more electron electron repulsion and a weaker attraction

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13
Q

molecular level adaptation of sodium chloride

A

-each Cl- is surrounded by 6 Na+ ions
- each Na+ is surrounded by 6 Cl- ions

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14
Q

how to expect a charge on an atom of a ceritan element

A

they want to get to the closest noble gas configuration

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15
Q

ion-ion interactions as electrostatic interactions

A
  • force of attraction between ions:
    F=k(q1q2 / r^2)
    -> more charge -> stronger attraction
    -> smaller ions (smaller r)-> stronger attraction
  • the stronger the force, the more energy is required to overcome it
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16
Q

Forming ion-ion interactions releases energy

A
  • energy released when an ionic lattice forms from ions in the gas phase (this is called lattice energy)
  • the amount of energy released depends on how strong the interaction between ions is
    – the stronger the attraction, the greater the lattice energy
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17
Q

Lattice energy and melting points

A

those ionic compounds with the highest lattice energy also have the highest melting point because there is a shorter distance resulting in the strongest force to be overcome

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18
Q

complex ionic compounds contain what

A

a polyatomic ion

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19
Q

polyatomic ion

A
  • a group of atoms covalently bonded to each other
  • the entire particle has a charge
    ex
    NH4+
    CN-
    NO3-
    CO3^2-
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20
Q

what is temperature

A
  • a useful macroscopic way of thinking about temperature is that it tells you in which direction thermal energy (often called heat) will move
  • energy always moves from a hotter object (higher temperature) to cooler (lower temperature) object
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21
Q

characteristics/conversions of temperature

A

-unites: celcius, Kalvin, Fahrenheit
-A change of 1 degree celsius = a change of 1 K
- 0 degrees celsius = 273.15 K
-0 K = absolute zero (lowest possible temp)

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22
Q

if you have one drop of boiling water vs a bucket of boiling water… do they have the same temperature?

A

yes, temperature depends on Average energy of particles

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23
Q

f you have one drop of boiling water vs a bucket of boiling water… do they have the same thermal energy?

A

no, thermal energy depends on total energy of particles, which depends on amount

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24
equation for temperature and kinetic energy of a monatomic gas
KEavg = 1/2 m (v avg)^2 T is directly related to the average kinetic energy of particles in a gas
25
at a given temperature, do all the atoms in a gas move at the same speed?
no
26
Can individual gas particles have a temperature?
no
27
populations vs individual molecules
- populations of molecules have a temperature where KEavg = 3/2 k-sub-b T KEavg = average kinetic energy in Joules k-sub-B = Boltzmann constant 138065 * 10^-23 J/K T = temperature K - one molecule can have a KE value, but NOT a temperature value -> temperature is a bulk property
28
Maxwell-Boltzmann Distribution for a gas at different temperatures
higer the avg velocity -> higher avg KE -> higher the T velocity is on x axis, meaning a higher average velocity is a curve spread further over the x axis
29
if samples of the same mass of different gases were taken at the same temperature, which would have the highest kinetic energy?
all average KE's would be the same
30
Ideal gas
consists of particles that have the following properties: -particles have no volume -- they are very small compared to the size of the container -particles exert no force on each other --particles are neither attracted to nor repelled from each other -kinetic energy is unchanged when gases collide with each other or the wall -ideal gases aren't real, but we can often model their behavior by a assuming a gas is ideal
31
describing ideal gases
through their macroscopic properties - pressure (P) - Volume (V) - temperature (T) - amount (moles) (n)
32
Boyle's law
V (is proportional to) 1/P - as V increases, P decreases - smaller V with higher P leads to more collisions with container (and same amount of moles) - larver V with lower P (and same amount of moles) leads to less collisions with container
33
charle's law
V (is proportional to) T as V increases T increases - tokeep P (pressure) constant (same # of collisions), particles need to move faster
34
Avogadro's law
As n increases, V increases
35
Ideal Gas law
summarizes the relationships between pressure, volume, temperature, and amount in a single equation - PV = nRT P= pressure (atmospheres) V = volume (Liters) n = moles R = constant (0.0821 L*atm / mol*K) T= temperature (Kalvins)
36
purpose of systems and surroundings
to figure out where the energy is coming from or where it is going - to keep track of energy changes
37
A system deff
part of the universe we are studying
38
the surroundings deff
Everything other than the system
39
open system def
can transfer both energy and matter in and out of system (eg biological systems)
40
closed system def
can transfer energy but not matter in and out of the system (ie, closed container)
41
isolated system def
can not transfer energy or matter in and out of the system (ie really good thermos)
42
Internal energy E (or U in physiscs)
- sum of all the KE and PE of all particles in the system - can't measure the exact value by can calculate the change : delta E = Efinal - Einitial
43
delta E notations
-energy out of the system to the surroundings: - sign (delta E < 0) -energy into system from surroundings: + sign (delta E >0)
44
motions of particles
described in terms of translation (across), rotation, and vibration
45
Thermal energy
molecular kinetic energy (translation + rotation + vibration) -> energy of random molecular motion - Units = Joules (J)
46
types of motion possible for molecules in each phase
gas: translation, rotation, vibration liquid: somewhat translation, somewhat rotation, vibration solid: vibration
47
heat
amount of energy that flows spontaneously from a hotter body to a cold body units: Joules (J)
48
temperature
a measure of the average kinetic energy of the molecules in a sample Units: Kalvins (K) KEavg = 3/2 RT
49
First law of thermodynamics
energy can not be created or destroyed (but can be transferred t=or transformed) delta E = q + w - q is the thermal energy (heat) change (J) - w is the work done (J) - in chemistry, most systems we study (except gases) do not expand or contract musch so we will ignore work; delta E = q
50
Endothermic process
heat goes into the system from the surroundings; q>0 (positive)
51
Exothermic process
heat goes from the system into the surroundings; q<0 (negative)
52
Enthalpy (delta H)
At a constant pressure, enthalpy (delta H) is the heat absorbed or emitted during a process
53
Enthalpy changes for phase changes
Ice to water is endothermic, water to vapor is endothermic Vapor to water is exothermic, water to ice is exothermic
54
Forming Hydrogen bonds is exothermic
- when water vapor condenses, heat is released - Hydrogen bonds form between molecules and decrease the chemical potential energy of the water molecules
55
Breaking Hydrogen bonds is endothermic
- When water boils, heat is absorbed - hydrogen bonds between molecules are broken and increase the chemical potential energy of the water molecules
56
is sweat evaporating or water freezing the endothermic process
sweat evaporating - water is absorbing energy from the surroundings (body) to evaporate
57
delta H of phase change
delta H melting = - delta H freezing delta H vaporizing = - delta H condensing
58
enthalpy change related to hydrogen bonds
the molecule with the most ability to create hydrogen bonds will require the most enthalpy
59
Specific heat capacity purpose and overall def
to quantify the heat from temperature change - how much energy it takes to change the temperature of a substance - both values are unique to each substance
60
specific heat
energy required to raise 1 g by 1 degree C
61
molar heat cpacity
energy required to raise 1 mol by 1 degree celcius
62
Heat and temperature related through heat capacity
the amount of heat (q) required to heat a substance by 1 degree C is directly related to: - amount of substance (m) -> more mass requires more heat to raise temp - identity of substance (related to strength of IMFs) -> higher heat capacity (c) requires more heat to raise temp q=mc deltaT
63
why does it take more energy to raise the temperature of liquid water than water vapor?
Heating steam: all the energy put in goes into increasing the average kinetic energy heating liquid H2O: some energy increases KEavg + some is used to overcome IMFs
64
How can we tell that the KEavg of a substance is increasing
- If the KEavg increases, the molecules are moving faster (translate, vibrate, rotate - we can't see how fast molecules are moving - what is our evidence that this is happening -> we can measure the temperature KEavg = 3/2 RT
65
Durring a phase change
- the temperature does not change -> it will start to change again after the phase change is complete - when melting or boiling, energy is added to the system from the surroundings and used to overcome the interactions between the particles - When condensing or freezing, energy from the formation of new interactions is being released into the surroundings
66

Chem 103 (4 decks)

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